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TABLE OF CONTENTS
Chapter 1
Chemical Foundations ......................................................................................... 1
Chapter 2
Atoms, Molecules, and Ions ...............................................................................25
Chapter 3
Stoichiometry .....................................................................................................46
Chapter 4
Types of Chemical Reactions and Solution Stoichiometry ..................................93
Chapter 5
Gases ...............................................................................................................139
Chapter 6
Thermochemistry .............................................................................................184
Chapter 7
Atomic Structure and Periodicity ......................................................................215
Chapter 8
Bonding: General Concepts .............................................................................250
Chapter 9
Covalent Bonding: Orbitals .............................................................................304
Chapter 10
Liquids and Solids ............................................................................................341
Chapter 11
Properties of Solutions .....................................................................................380
Chapter 12
Chemical Kinetics ............................................................................................418
Chapter 13
Chemical Equilibrium ......................................................................................458
Chapter 14
Acids and Bases ...............................................................................................500
Chapter 15
Acid-Base Equilibria ........................................................................................563
Chapter 16
Solubility and Complex Ion Equilibria .............................................................621
Chapter 17
Spontaneity, Entropy, and Free Energy.............................................................659
Chapter 18
Electrochemistry ..............................................................................................688
Chapter 19
The Nucleus: A Chemist’s View ......................................................................741
Chapter 20
The Representative Elements............................................................................760
Chapter 21
Transition Metals and Coordination Chemistry.................................................782
Chapter 22
Organic and Biological Molecules ....................................................................810
iii
CHAPTER 1
CHEMICAL FOUNDATIONS
Questions
17.
A law summarizes what happens, e.g., law of conservation of mass in a chemical reaction or
the ideal gas law, PV = nRT. A theory (model) is an attempt to explain why something
happens. Dalton’s atomic theory explains why mass is conserved in a chemical reaction. The
kinetic molecular theory explains why pressure and volume are inversely related at constant
temperature and moles of gas present, as well as explaining the other mathematical
relationships summarized in PV = nRT.
18.
A dynamic process is one that is active as opposed to static. In terms of the scientific
method, scientists are always performing experiments to prove or disprove a hypothesis or a
law or a theory. Scientists do not stop asking questions just because a given theory seems to
account satisfactorily for some aspect of natural behavior. The key to the scientific method is
to continually ask questions and perform experiments. Science is an active process, not a
static one.
19.
The fundamental steps are
(1) making observations;
(2) formulating hypotheses;
(3) performing experiments to test the hypotheses.
The key to the scientific method is performing experiments to test hypotheses. If after the test
of time the hypotheses seem to account satisfactorily for some aspect of natural behavior,
then the set of tested hypotheses turns into a theory (model). However, scientists continue to
perform experiments to refine or replace existing theories.
20.
A random error has equal probability of being too high or too low. This type of error occurs
when estimating the value of the last digit of a measurement. A systematic error is one that
always occurs in the same direction, either too high or too low. For example, this type of
error would occur if the balance you were using weighed all objects 0.20 g too high, that is, if
the balance wasn’t calibrated correctly. A random error is an indeterminate error, whereas a
systematic error is a determinate error.
21.
A qualitative observation expresses what makes something what it is; it does not involve a
number; e.g., the air we breathe is a mixture of gases, ice is less dense than water, rotten milk
stinks.
The SI units are mass in kilograms, length in meters, and volume in the derived units of m3.
The assumed uncertainty in a number is ±1 in the last significant figure of the number. The
precision of an instrument is related to the number of significant figures associated with an
1
2
CHAPTER 1
CHEMICAL FOUNDATIONS
experimental reading on that instrument. Different instruments for measuring mass, length, or
volume have varying degrees of precision. Some instruments only give a few significant
figures for a measurement, whereas others will give more significant figures.
22.
Precision: reproducibility; accuracy: the agreement of a measurement with the true value.
a. Imprecise and inaccurate data: 12.32 cm, 9.63 cm, 11.98 cm, 13.34 cm
b. Precise but inaccurate data: 8.76 cm, 8.79 cm, 8.72 cm, 8.75 cm
c. Precise and accurate data: 10.60 cm, 10.65 cm, 10.63 cm, 10.64 cm
Data can be imprecise if the measuring device is imprecise as well as if the user of the
measuring device has poor skills. Data can be inaccurate due to a systematic error in the
measuring device or with the user. For example, a balance may read all masses as weighing
0.2500 g too high or the user of a graduated cylinder may read all measurements 0.05 mL too
low.
A set of measurements that are imprecise implies that all the numbers are not close to each
other. If the numbers aren’t reproducible, then all the numbers can’t be very close to the true
value. Some say that if the average of imprecise data gives the true value, then the data are
accurate; a better description is that the data takers are extremely lucky.
23.
Significant figures are the digits we associate with a number. They contain all of the certain
digits and the first uncertain digit (the first estimated digit). What follows is one thousand
indicated to varying numbers of significant figures: 1000 or 1 × 103 (1 S.F.); 1.0 × 103 (2
S.F.); 1.00 × 103 (3 S.F.); 1000. or 1.000 × 103 (4 S.F.).
To perform the calculation, the addition/subtraction significant figure rule is applied to 1.5 −
1.0. The result of this is the one-significant-figure answer of 0.5. Next, the multiplication/division rule is applied to 0.5/0.50. A one-significant-figure number divided by a
two-significant-figure number yields an answer with one significant figure (answer = 1).
24.
From Figure 1.9 of the text, a change in temperature of 180°F is equal to a change in
temperature of 100°C and 100 K. A degree unit on the Fahrenheit scale is not a large as a
degree unit on the Celsius or Kelvin scales. Therefore, a 20° change in the Celsius or Kelvin
temperature would correspond to a larger temperature change than a 20° change in the
Fahrenheit scale. The 20° temperature change on the Celsius and Kelvin scales are equal to
each other.
25.
Straight line equation: y = mx + b, where m is the slope of the line and b is the y-intercept. For
the TF vs. TC plot:
TF = (9/5)TC + 32
y= m x + b
The slope of the plot is 1.8 (= 9/5) and the y-intercept is 32°F.
For the TC vs. TK plot:
TC = TK − 273
y= mx + b
The slope of the plot is 1, and the y-intercept is −273°C.
CHAPTER 1
26.
CHEMICAL FOUNDATIONS
3
a. coffee; saltwater; the air we breathe (N2 + O2 + others); brass (Cu + Zn)
b. book; human being; tree; desk
c. sodium chloride (NaCl); water (H2O); glucose (C6H12O6); carbon dioxide (CO2)
d. nitrogen (N2); oxygen (O2); copper (Cu); zinc (Zn)
e. boiling water; freezing water; melting a popsicle; dry ice subliming
f.
Elecrolysis of molten sodium chloride to produce sodium and chlorine gas; the explosive
reaction between oxygen and hydrogen to produce water; photosynthesis, which converts
H2O and CO2 into C6H12O6 and O2; the combustion of gasoline in our car to produce CO2
and H2O
Exercises
Significant Figures and Unit Conversions
27.
a.
exact
c. exact
28.
29.
30.
31.
b. inexact
d. inexact (π has an infinite number of decimal places.)
a. one significant figure (S.F.). The implied uncertainty is ±1000 pages. More significant
figures should be added if a more precise number is known.
b. two S.F.
c. four S.F.
d. two S.F.
e. infinite number of S.F. (exact number)
a. 6.07 × 10 −15 ; 3 S.F.
b. 0.003840; 4 S.F.
c. 17.00; 4 S.F.
d. 8 × 108; 1 S.F.
e. 463.8052; 7 S.F.
f.
g. 301; 3 S.F.
h. 300.; 3 S.F.
a. 100; 1 S.F.
b. 1.0 × 102; 2 S.F.
c. 1.00 × 103; 3 S.F.
d. 100.; 3 S.F.
e. 0.0048; 2 S.F.
f.
g. 4.80 × 10 −3 ; 3 S.F.
h. 4.800 × 10 −3 ; 4 S.F.
f.
one S.F.
300; 1 S.F.
0.00480; 3 S.F.
When rounding, the last significant figure stays the same if the number after this significant
figure is less than 5 and increases by one if the number is greater than or equal to 5.
a. 3.42 × 10 −4
b. 1.034 × 104
c. 1.7992 × 101
d. 3.37 × 105
32.
a. 4 × 105
b. 3.9 × 105
c. 3.86 × 105
d. 3.8550 × 105
33.
Volume measurements are estimated to one place past the markings on the glassware. The
first graduated cylinder is labeled to 0.2 mL volume increments, so we estimate volumes to
4
CHAPTER 1
CHEMICAL FOUNDATIONS
the hundredths place. Realistically, the uncertainty in this graduated cylinder is ±0.05 mL.
The second cylinder, with 0.02 mL volume increments, will have an uncertainty of ±0.005
mL. The approximate volume in the first graduated cylinder is 2.85 mL, and the volume in
the other graduated cylinder is approximately 0.280 mL. The total volume would be:
2.85 mL
+0.280 mL
3.13 mL
We should report the total volume to the hundredths place because the volume from the first
graduated cylinder is only read to the hundredths (read to two decimal places). The first
graduated cylinder is the least precise volume measurement because the uncertainty of this
instrument is in the hundredths place, while the uncertainty of the second graduated cylinder
is to the thousandths place. It is always the lease precise measurement that limits the
precision of a calculation.
34.
a. Volumes are always estimated to one position past the marked volume increments. The
estimated volume of the first beaker is 32.7 mL, the estimated volume of the middle
beaker is 33 mL, and the estimated volume in the last beaker is 32.73 mL.
b. Yes, all volumes could be identical to each other because the more precise volume
readings can be rounded to the other volume readings. But because the volumes are in
three different measuring devices, each with its own unique uncertainty, we cannot say
with certainty that all three beakers contain the same amount of water.
c. 32.7 mL
33 mL
32.73 mL
98.43 mL = 98 mL
The volume in the middle beaker can only be estimated to the ones place, which dictates that
the sum of the volume should be reported to the ones place. As is always the case, the least
precise measurement determines the precision of a calculation.
35.
For addition and/or subtraction, the result has the same number of decimal places as the
number in the calculation with the fewest decimal places. When the result is rounded to the
correct number of significant figures, the last significant figure stays the same if the number
after this significant figure is less than 5 and increases by one if the number is greater than or
equal to 5. The underline shows the last significant figure in the intermediate answers.
a. 212.2 + 26.7 + 402.09 = 640.99 = 641.0
b. 1.0028 + 0.221 + 0.10337 = 1.32717 = 1.327
c. 52.331 + 26.01 − 0.9981 = 77.3429 = 77.34
d. 2.01 × 102 + 3.014 × 103 = 2.01 × 102 + 30.14 × 102 = 32.15 × 102 = 3215
When the exponents are different, it is easiest to apply the addition/subtraction rule when
all numbers are based on the same power of 10.
e. 7.255 − 6.8350 = 0.42 = 0.420 (first uncertain digit is in the third decimal place).
CHAPTER 1
36.
CHEMICAL FOUNDATIONS
5
For multiplication and/or division, the result has the same number of significant figures as the
number in the calculation with the fewest significant figures.
a.
0.102 × 0.0821 × 273
= 2.2635 = 2.26
1.01
b. 0.14 × 6.022 × 1023 = 8.431 × 1022 = 8.4 × 1022; since 0.14 only has two significant
figures, the result should only have two significant figures.
c. 4.0 × 104 × 5.021 × 10 −3 × 7.34993 × 102 = 1.476 × 105 = 1.5 × 105
d.
37.
2.00 × 106
3.00 × 10
−7
= 6.6667 × 1012 = 6.67 × 1012
a. Here, apply the multiplication/division rule first; then apply the addition/subtraction rule
to arrive at the one-decimal-place answer. We will generally round off at intermediate
steps in order to show the correct number of significant figures. However, you should
round off at the end of all the mathematical operations in order to avoid round-off error.
The best way to do calculations is to keep track of the correct number of significant
figures during intermediate steps, but round off at the end. For this problem, we
underlined the last significant figure in the intermediate steps.
2.526 0.470
80.705
= 0.8148 + 0.7544 + 186.558 = 188.1
+
+
3.1
0.623
0.4326
b. Here, the mathematical operation requires that we apply the addition/subtraction rule
first, then apply the multiplication/division rule.
6.404 × 2.91 6.404 × 2.91
=
= 12
18.7 − 17.1
1.6
c. 6.071 × 10 −5 − 8.2 × 10 −6 − 0.521 × 10 −4 = 60.71 × 10 −6 − 8.2 × 10 −6 − 52.1 × 10 −6
= 0.41 × 10 −6 = 4 × 10 −7
d.
3.8 × 10 −12 + 4.0 × 10 −13
38 × 10 −13 + 4.0 × 10 −13
42 × 10 −13
=
=
= 6.3 × 10 − 26
12
13
12
12
12
4 × 10 + 6.3 × 10
4 × 10 + 63 × 10
67 × 10
e.
9.5 + 4.1 + 2.8 + 3.175 19.575
= 4.89 = 4.9
=
4
4
Uncertainty appears in the first decimal place. The average of several numbers can only
be as precise as the least precise number. Averages can be exceptions to the significant
figure rules.
f.
38.
8.925 − 8.905
0.020
× 100 = 0.22
× 100 =
8.925
8.925
a. 6.022 × 1023 × 1.05 × 102 = 6.32 × 1025
6
CHAPTER 1
b.
6.6262 × 10 −34 × 2.998 × 108
2.54 × 10 −9
CHEMICAL FOUNDATIONS
= 7.82 × 10 −17
c. 1.285 × 10 −2 + 1.24 × 10 −3 + 1.879 × 10 −1
= 0.1285 × 10 −1 + 0.0124 × 10 −1 + 1.879 × 10 −1 = 2.020 × 10 −1
When the exponents are different, it is easiest to apply the addition/subtraction rule when
all numbers are based on the same power of 10.
d.
e.
f.
39.
(1.00866 − 1.00728)
6.02205 × 10
23
9.875 × 10
6.02205 × 10
2
× 100 =
23
= 2.29 × 10 −27
0.080 × 10 2
9.875 × 10
2
× 100 = 8.1 × 10 −1
9.42 × 10 2 + 8.234 × 10 2 + 1.625 × 103
0.942 × 103 + 0.824 × 103 + 1.625 × 103
=
3
3
= 1.130 × 103
a. 8.43 cm ×
1000 mm
1m
= 84.3 mm
×
m
100 cm
1m
1 × 10 nm
9
f.
903.3 nm ×
a.
1 Tg ×
1m
1 × 109 nm
×
b. 2.41 × 102 cm ×
e. 235.3 m ×
1000 mm
= 2.353 × 105 mm
m
1 × 106 μm
= 0.9033 μm
m
1 × 1012 m 1 × 109 nm
= 6.50 × 10 23 nm
×
Tm
m
1g
c. 25 fg ×
1 × 10 fg
15
d. 8.0 dm3 ×
e. 1 mL ×
1L
dm
3
×
1 kg
= 25 × 10 −18 kg = 2.5 × 10 −17 kg
1000 g
= 8.0 L (1 L = 1 dm3 = 1000 cm3 = 1000 mL)
1L
1 × 10 6 μL
×
= 1 × 103 μL
1 000 mL
L
1g
1 × 10 6 μg
×
1m
= 2.41 m
100 cm
100 cm
= 2.945 × 10 −5 cm
m
1 × 1012 g
1 kg
= 1 × 109 kg
×
Tg
1000 g
b. 6.50 × 102 Tm ×
1 μg ×
×
1 km
= 14.45 km
1000 m
d. 1.445 × 104 m ×
f.
0.00138
9.875 × 10 2 − 9.795 × 10 2
c. 294.5 nm ×
40.
=
1 × 1012 pg
= 1 × 106 pg
g
CHAPTER 1
41.
CHEMICAL FOUNDATIONS
7
a. Appropriate conversion factors are found in Appendix 6. In general, the number of
significant figures we use in the conversion factors will be one more than the number of
significant figures from the numbers given in the problem. This is usually sufficient to
avoid round-off error.
1 lb
16 oz
3.91 kg ×
= 8.62 lb; 0.62 lb ×
= 9.9 oz
0.4536 kg
lb
Baby’s weight = 8 lb and 9.9 oz or, to the nearest ounce, 8 lb and 10. oz.
51.4 cm ×
1 in
= 20.2 in ≈ 20 1/4 in = baby’s height
2.54 cm
b. 25,000 mi ×
1000 m
1.61 km
= 4.0 × 104 km; 4.0 × 104 km ×
= 4.0 × 107 m
km
mi

1m  
1m 
 ×  2.1 dm ×
 = 1.2 × 10 −2 m3
c. V = 1 × w × h = 1.0 m ×  5.6 cm ×
100 cm  
10 dm 

3
1L
 1 0 dm 
1.2 × 10 −2 m3 × 
= 12 L
 ×
dm 3
 m 
3
12 L ×
42.
43.
1000 cm 3  1 in 
 = 730 in3; 730 in3 ×
× 
L
2
.
54
cm


a. 908 oz ×
1 lb
0.4536 kg
×
= 25.7 kg
16 oz
lb
b. 12.8 L ×
1 qt
1 gal
×
= 3.38 gal
0.9463 L
4 qt
c. 125 mL ×
1L
1 qt
×
= 0.132 qt
1000 mL 0.9463 L
d. 2.89 gal ×
4 qt
1L
1000 mL
×
×
= 1.09 × 104 mL
1 gal 1 .057 qt
1L
e. 4.48 lb ×
453.6 g
= 2.03 × 103 g
1 lb
f.
1L
1.06 qt
×
= 0.58 qt
1000 mL
L
550 mL ×
a. 1.25 mi ×
3
 1 ft 

 = 0.42 ft3
12
in


40 rods
8 furlongs
= 10.0 furlongs; 10.0 furlongs ×
= 4.00 × 102 rods
furlong
mi
4.00 × 102 rods ×
5.5 yd 36 in 2.54 cm
1m
×
×
×
= 2.01 × 103 m
rod
yd
in
100 cm
8
CHAPTER 1
2.01 × 103 m ×
CHEMICAL FOUNDATIONS
1 km
= 2.01 km
1000 m
b. Let's assume we know this distance to ±1 yard. First, convert 26 miles to yards.
26 mi ×
5280 ft 1 yd
×
= 45,760. yd
mi
3 ft
26 mi + 385 yd = 45,760. yd + 385 yd = 46,145 yards
46,145 yard ×
46,145 yard ×
1 rod
1 furlong
= 8390.0 rods; 8390.0 rods ×
= 209.75 furlongs
5.5 yd
40 rods
1 km
36 in 2.54 cm
1m
×
×
= 42,195 m; 42,195 m ×
yd
in
100 cm
1000 m
= 42.195 km
2
 1 km 
10,000 m
 = 1 × 10 −2 km 2
× 
ha
 1000 m 
2
44.
a. 1 ha ×
b. 5.5 acre ×
160 rod 2  5.5 yd 36 in 2.54 cm
1m
× 
×
×
×
acre
yd
in
100 cm
 rod
4
2
2.2 × 10 m ×
1 ha
1 × 10 4 m 2
2

 = 2.2 × 104 m2

2
 1 km 
 = 0.022 km2
= 2.2 ha; 2.2 × 10 m × 
 1000 m 
4
2
c. Area of lot = 120 ft × 75 ft = 9.0 × 103 ft2
 1 yd
1 rod
9.0 × 10 ft × 
×
5.5 yd
 3 ft
3
2
2

$6,500
$31,000
1 acre
 ×
=
= 0.21 acre;
2
0.21 acre
acre
160 rod

We can use our result from (b) to get the conversion factor between acres and hectares
(5.5 acre = 2.2 ha.). Thus 1 ha = 2.5 acre.
1 ha
$6,500
$77,000
=
= 0.084 ha; the price is:
2.5 acre
0.084 ha
ha
0.21 acre ×
45.
a. 1 troy lb ×
12 troy oz
20 pw
24 grains 0.0648 g
1 kg
×
×
×
×
= 0.373 kg
troy lb
troy oz
pw
grain
1000 g
1 troy lb = 0.373 kg ×
b. 1 troy oz ×
2.205 lb
= 0.822 lb
kg
20 pw
24 grains 0.0648 g
×
×
= 31.1 g
troy oz
pw
grain
1 troy oz = 31.1 g ×
1 carat
= 156 carats
0.200 g
CHAPTER 1
CHEMICAL FOUNDATIONS
c. 1 troy lb = 0.373 kg; 0.373 kg ×
46.
a. 1 grain ap ×
9
1000 g 1 cm 3
= 19.3 cm3
×
kg
19.3 g
1 scruple
1 dram ap
3.888 g
×
×
= 0.06480 g
20 grain ap
3 scruples
dram ap
From the previous question, we are given that 1 grain troy = 0.0648 g = 1 grain ap. So the
two are the same.
b. 1 oz ap ×
8 dram ap
3.888 g
1 oz troy *
×
×
= 1.00 oz troy; *see Exercise 45b.
oz ap
dram ap
31.1 g
c. 5.00 × 102 mg ×
0.386 scruple ×
d. 1 scruple ×
1g
1 dram ap
3 scruples
×
×
= 0.386 scruple
1000 mg
3.888 g
dram ap
20 grains ap
= 7.72 grains ap
scruple
1 dram ap
3.888 g
×
= 1.296 g
3 scruples
dram ap
1 capsule
= 24 capsules
0.65 g
47.
15.6 g ×
48.
1.5 teaspoons ×
80. mg acet
= 240 mg acetaminophen
0.50 teaspoon
240 mg acet
1 lb
×
= 22 mg acetaminophen/kg
24 lb
0.454 kg
240 mg acet
1 lb
×
= 15 mg acetaminophen/kg
35 lb
0.454 kg
The range is from 15 to 22 mg acetaminophen per kg of body weight.
49.

3.00 × 10 8 m  1.094 yd
60 s
60 min
1 knot
 ×
warp 1.71 =  5.00 ×
×
×
×

s
m
min
h
2030 yd/h


= 2.91 × 109 knots
8


1 mi
60 s
60 min
 5.00 × 3.00 × 10 m  × 1 km ×
= 3.36 × 109 mi/h
×
×


s
1.609 km
min
h

 1000 m
50.
100. m
100. m
1 km
60 s
60 min
×
×
×
= 10.4 m/s;
= 37.6 km/h
9.58 s
9.58 s 1000 m
min
h
10
CHAPTER 1
CHEMICAL FOUNDATIONS
100. m 1.0936 yd 3 ft
34.2 ft
1 mi
60 s
60 min
×
×
×
×
×
= 34.2 ft/s;
= 23.3 mi/h
9.58 s
m
yd
s
5280 ft
min
h
1.00 × 102 yd ×
51.
1m
9.58 s
×
= 8.76 s
1.0936 yd
100. m
65 km 0.6214 mi
= 40.4 = 40. mi/h
×
h
km
To the correct number of significant figures (2), 65 km/h does not violate a 40 mi/h speed
limit.
52.
53.
112 km ×
0.6214 mi
1h
×
= 1.1 h = 1 h and 6 min
km
65 mi
112 km ×
0.6214 mi
1 gal
3.785 L
×
×
= 9.4 L of gasoline
km
28 mi
gal
$1.32
1 kg
2.45 euros
×
×
= $1.47/lb
euro
2.2046 lb
kg
One pound of peaches costs $1.47.
54.
For the gasoline car:
500. mi ×
1 gal
$3.50
×
28.0 mi
gal
= $62.5
For the E85 car:
500. mi ×
1 gal
$2.85
×
= $63.3
22.5 mi
gal
The E85 vehicle would cost slightly more to drive 500. miles as compared to the gasoline
vehicle ($63.3 versus $62.5).
2
55.
 5280 ft 
10 3
Volume of lake = 100 mi2 × 
 × 20 ft = 6 × 10 ft
mi


3
2.54 cm 
1 mL 0.4 μg
 12 in
= 7 × 1014 μg mercury
6 × 1010 ft3 × 
×
×
 ×
3
in 
mL
cm
 ft
7 × 1014 μg ×
56.
1g
1 × 10 μg
6
×
1 kg
1 × 103 g
= 7 × 105 kg of mercury
Volume of room = 18 ft × 12 ft × 8 ft = 1700 ft3 (carrying one extra significant figure)
CHAPTER 1
CHEMICAL FOUNDATIONS
3
11
3
3
 1m 
 12 in 
 2.54 cm 
 = 48 m 3
1700 ft × 
 × 
 × 
 ft 
 in 
 100 cm 
3
48 m3 ×
400,000 μg CO
m
3
×
1 g CO
1 × 10 6 μg CO
= 19 g = 20 g CO (to 1 sig. fig.)
Temperature
57.
a. TC =
b. TC =
c. TC =
d. TC =
58.
5
9
5
9
5
9
5
9
(TF − 32) =
5
9
(−459°F − 32) = −273°C; TK = TC + 273 = −273°C + 273 = 0 K
(−40.°F − 32) = −40.°C; TK = −40.°C + 273 = 233 K
(68°F − 32) = 20.°C; TK = 20.°C + 273 = 293 K
(7 × 107°F − 32) = 4 × 107°C; TK = 4 × 107°C + 273 = 4 × 107 K
96.1°F ±0.2°F; first, convert 96.1°F to °C. TC =
5
5
(TF − 32) = (96.1 − 32) = 35.6°C
9
9
A change in temperature of 9°F is equal to a change in temperature of 5°C.
uncertainty is:
5° C
= ±0.1°C. Thus 96.1 ±0.2°F = 35.6 ±0.1°C.
9° F
9
9
a. TF =
× TC + 32 =
× 39.2°C + 32 = 102.6°F (Note: 32 is exact.)
5
5
±0.2°F ×
59.
TK = TC + 273.2 = 39.2 + 273.2 = 312.4 K
b. TF =
c.
TF =
d. TF =
60.
9
× (−25) + 32 = −13°F; TK = −25 + 273 = 248 K
5
9
× (−273) + 32 = −459°F; TK = −273 + 273 = 0 K
5
9
× 801 + 32 = 1470°F; TK = 801 + 273 = 1074 K
5
a. TC = TK − 273 = 233 − 273 = -40.°C
TF =
9
9
× TC + 32 =
× (−40.) + 32 = −40.°F
5
5
b. TC = 4 − 273 = −269°C; TF =
9
× (−269) + 32 = −452°F
5
So the
12
CHAPTER 1
c. TC = 298 − 273 = 25°C; TF =
9
× 25 + 32 = 77°F
5
d. TC = 3680 − 273 = 3410°C; TF =
61.
TF =
CHEMICAL FOUNDATIONS
9
× 3410 + 32 = 6170°F
5
9
× TC + 32; from the problem, we want the temperature where TF = 2TC.
5
Substituting:
2TC =
9
32
× TC + 32, (0.2)TC = 32, TC =
= 160°C
5
0.2
TF = 2TC when the temperature in Fahrenheit is 2(160) = 320°F. Because all numbers when
solving the equation are exact numbers, the calculated temperatures are also exact numbers.
62.
TC =
5
5
(TF – 32) = (72 – 32) = 22°C
9
9
TC = TK – 273 = 313 – 273 = 40.°C
The difference in temperature between Jupiter at 313 K and Earth at 72°F is 40.°C – 22 °C =
18°C.
63.
a. A change in temperature of 140°C is equal to 50°X. Therefore,
140 o C
is the unit con50 o X
version between a degree on the X scale to a degree on the Celsius scale. To account for
the different zero points, −10° must be subtracted from the temperature on the X scale to
get to the Celsius scale. The conversion between °X to °C is:
TC = TX ×
140 o C
50 o X
− 10°C, TC = TX ×
14 o C
5o X
− 10°C
The conversion between °C to °X would be:
TX = (TC + 10°C)
b. Assuming 10°C and
5o X
14 o C
5o X
14 o C
TX = (22.0°C + 10°C)
are exact numbers:
5o X
14 o C
= 11.4°X
c. Assuming exact numbers in the temperature conversion formulas:
TC = 58.0°X ×
14 o C
5o X
− 10°C = 152°C
CHAPTER 1
CHEMICAL FOUNDATIONS
13
TK = 152°C + 273 = 425 K
TF =
64.
9o F
5o C
a.
× 152°C + 32°F = 306°F
100oA
115oC
100oA
160oC
0oA
-45oC
A change in temperature of 160°C equals a
change in temperature of 100°A.
160°C
is our unit conversion for a
100°A
degree change in temperature.
So
At the freezing point: 0°A = −45°C
Combining these two pieces of information:
TA = (TC + 45°C) ×
b. TC = (TF − 32) ×
TF − 32 =
c. TC = TA ×
8°C
100°A
5°A
= (TC + 45°C) ×
or TC = TA ×
− 45°C
5°A
160°C
8°C
8
5
5
; TC = TA ×
− 45 = (TF − 32) ×
9
5
9
8
72
72°F
9


×  TA × − 45  = TA ×
− 81, TF = TA ×
− 49°F
5
25
25°A
5


3Tc
8
8
− 45 and TC = TA; so TC = TC ×
− 45,
= 45, TC = 75°C = 75°A
5
5
5
d. TC = 86°A ×
8°C
72°F
− 45°C = 93°C; TF = 86°A ×
− 49°F = 199°F = 2.0 × 102°F
5°A
25°A
e. TA = (45°C + 45°C) ×
5°A
= 56°A
8°C
Density
3
65.
453.6 g
 2.54 cm 
5
3
= 1.6 × 105 g; V = 1.2 × 104 in3 × 
 = 2.0 × 10 cm
in
lb


5
mass
1 × 10 g
Density =
=
= 0.80 g/cm 3
5
3
volume
2.0 × 10 cm
Mass = 350 lb ×
Because the material has a density less than water, it will float in water.
66.
V =
4 3 4
2.0 g
= 3.8 g/cm3
π r = × 3.14 × (0.50 cm)3 = 0.52 cm 3 ; d =
3
3
3
0.52 cm
The ball will sink.
14
67.
CHAPTER 1
V=
3
4 3
4
1000 m 100 cm 

33
3
π r = × 3.14 ×  7.0 × 105 km ×
×
 = 1.4 × 10 cm
3
3
km
m 

Density =
68.
mass
=
volume
615.0 g
1.0 × 10 cm
2
3
=
6.2 g
cm 3
0.200 g 1 cm 3
= 0.28 cm3
×
carat
3.51 g
a. 5.0 carat ×
b. 2.8 mL ×
70.
1000 g
kg
= 1.4 × 106 g/cm3 = 1 × 106 g/cm3
33
3
1.4 × 10 cm
2 × 1036 kg ×
V = l × w × h = 2.9 cm × 3.5 cm × 10.0 cm = 1.0 × 102 cm3
d = density =
69.
CHEMICAL FOUNDATIONS
1 cm 3
3.51 g
1 carat
= 49 carats
×
×
3
mL
0.200 g
cm
For ethanol: 100. mL ×
0.789 g
= 78.9 g
mL
For benzene: 1.00 L ×
1000 mL 0.880 g
= 880. g
×
L
mL
Total mass = 78.9 g + 880. g = 959 g
71.
V = 21.6 mL − 12.7 mL = 8.9 mL; density =
72.
5.25 g ×
1 cm 3
10.5 g
33.42 g
= 3.8 g/mL = 3.8 g/cm3
8.9 mL
= 0.500 cm3 = 0.500 mL
The volume in the cylinder will rise to 11.7 mL (11.2 mL + 0.500 mL = 11.7 mL).
73.
a. Both have the same mass of 1.0 kg.
b. 1.0 mL of mercury; mercury is more dense than water. Note: 1 mL = 1 cm3.
13.6 g
0.998 g
= 14 g of mercury; 1.0 mL ×
= 1.0 g of water
mL
mL
1.0 mL ×
c. Same; both represent 19.3 g of substance.
19.3 mL ×
0.9982 g
19.32 g
= 19.3 g of water; 1.00 mL ×
= 19.3 g of gold
mL
mL
d. 1.0 L of benzene (880 g versus 670 g)
75 mL ×
8.96 g
1000 mL 0.880 g
= 670 g of copper; 1.0 L ×
= 880 g of benzene
×
mL
L
mL
CHAPTER 1
74.
CHEMICAL FOUNDATIONS
a. 1.50 qt ×
15
1L
1000 mL 0.789 g
×
×
= 1120 g ethanol
1.0567 qt
L
mL
3
13.6 g
 2.54 cm 
= 780 g mercury
b. 3.5 in3 × 
 ×
cm 3
 in 
75.
a. 1.0 kg feather; feathers are less dense than lead.
b. 100 g water; water is less dense than gold.
76.
1 cm 3
= 3.0 × 105 cm3 [H2(g) = hydrogen gas.]
0.000084 g
a. H2(g): V = 25.0 g ×
b. H2O(l): V = 25.0 g ×
c. Fe(s): V = 25.0 g ×
c. Same; both volumes are 1.0 L.
1 cm 3
= 25.0 cm3 [H2O(l) = water.]
0.9982 g
1 cm 3
= 3.18 cm3 [Fe(s) = iron.]
7.87 g
Notice the huge volume of the gaseous H2 sample as compared to the liquid and solid
samples. The same mass of gas occupies a volume that is over 10,000 times larger than the
liquid sample. Gases are indeed mostly empty space.
77.
V = 1.00 × 103 g ×
1 cm 3
= 44.3 cm3
22.57 g
44.3 cm3 = 1 × w × h = 4.00 cm × 4.00 cm × h, h = 2.77 cm
78.
V = 22 g ×
1 cm 3
= 2.5 cm3; V = πr2 × l, where l = length of the wire
8.96 g
2
 1 cm 
 0.25 mm 
 × l, l = 5.1 × 103 cm = 170 ft
2.5 cm = π × 
 × 
2


 10 mm 
2
3
Classification and Separation of Matter
79.
A gas has molecules that are very far apart from each other, whereas a solid or liquid has
molecules that are very close together. An element has the same type of atom, whereas a
compound contains two or more different elements. Picture i represents an element that
exists as two atoms bonded together (like H2 or O2 or N2). Picture iv represents a compound
(like CO, NO, or HF). Pictures iii and iv contain representations of elements that exist as
individual atoms (like Ar, Ne, or He).
a. Picture iv represents a gaseous compound. Note that pictures ii and iii also contain a
gaseous compound, but they also both have a gaseous element present.
b. Picture vi represents a mixture of two gaseous elements.
16
CHAPTER 1
CHEMICAL FOUNDATIONS
c. Picture v represents a solid element.
d. Pictures ii and iii both represent a mixture of a gaseous element and a gaseous compound.
80.
Solid: rigid; has a fixed volume and shape; slightly compressible
Liquid: definite volume but no specific shape; assumes shape of the container; slightly
Compressible
Gas: no fixed volume or shape; easily compressible
Pure substance: has constant composition; can be composed of either compounds or elements
Element: substances that cannot be decomposed into simpler substances by chemical or
physical means.
Compound: a substance that can be broken down into simpler substances (elements) by
chemical processes.
Homogeneous mixture: a mixture of pure substances that has visibly indistinguishable parts.
Heterogeneous mixture: a mixture of pure substances that has visibly distinguishable parts.
Solution: a homogeneous mixture; can be a solid, liquid or gas
Chemical change: a given substance becomes a new substance or substances with different
properties and different composition.
Physical change: changes the form (g, l, or s) of a substance but does no change the chemical
composition of the substance.
81.
Homogeneous: Having visibly indistinguishable parts (the same throughout).
Heterogeneous: Having visibly distinguishable parts (not uniform throughout).
a. heterogeneous (due to hinges, handles, locks, etc.)
b. homogeneous (hopefully; if you live in a heavily polluted area, air may be
heterogeneous.)
82.
c. homogeneous
d. homogeneous (hopefully, if not polluted)
e. heterogeneous
f.
a. heterogeneous
b. homogeneous
c. heterogeneous
d. homogeneous (assuming no imperfections in the glass)
heterogeneous
e. heterogeneous (has visibly distinguishable parts)
CHAPTER 1
83.
CHEMICAL FOUNDATIONS
a.
pure
b.
mixture
f.
pure
g. mixture
c.
17
mixture
d.
pure
h. mixture
i.
mixture
e. mixture (copper and zinc)
Iron and uranium are elements. Water (H2O) is a compound because it is made up of two or
more different elements. Table salt is usually a homogeneous mixture composed mostly of
sodium chloride (NaCl), but will usually contain other substances that help absorb water
vapor (an anticaking agent).
84.
Initially, a mixture is present. The magnesium and sulfur have only been placed together in
the same container at this point, but no reaction has occurred. When heated, a reaction occurs.
Assuming the magnesium and sulfur had been measured out in exactly the correct ratio for
complete reaction, the remains after heating would be a pure compound composed of
magnesium and sulfur. However, if there were an excess of either magnesium or sulfur, the
remains after reaction would be a mixture of the compound produced and the excess reactant.
85.
Chalk is a compound because it loses mass when heated and appears to change into another
substance with different physical properties (the hard chalk turns into a crumbly substance).
86.
Because vaporized water is still the same substance as solid water (H2O), no chemical
reaction has occurred. Sublimation is a physical change.
87.
A physical change is a change in the state of a substance (solid, liquid, and gas are the three
states of matter); a physical change does not change the chemical composition of the
substance. A chemical change is a change in which a given substance is converted into
another substance having a different formula (composition).
a. Vaporization refers to a liquid converting to a gas, so this is a physical change. The
formula (composition) of the moth ball does not change.
b. This is a chemical change since hydrofluoric acid (HF) is reacting with glass (SiO2) to
form new compounds that wash away.
c. This is a physical change because all that is happening during the boiling process is the
conversion of liquid alcohol to gaseous alcohol. The alcohol formula (C2H5OH) does not
change.
d. This is a chemical change since the acid is reacting with cotton to form new compounds.
88.
a. Distillation separates components of a mixture, so the orange liquid is a mixture (has an
average color of the yellow liquid and the red solid). Distillation utilizes boiling point
differences to separate out the components of a mixture. Distillation is a physical change
because the components of the mixture do not become different compounds or elements.
b. Decomposition is a type of chemical reaction. The crystalline solid is a compound, and
decomposition is a chemical change where new substances are formed.
c. Tea is a mixture of tea compounds dissolved in water. The process of mixing sugar into
tea is a physical change. Sugar doesn’t react with the tea compounds, it just makes the
solution sweeter.
18
CHAPTER 1
CHEMICAL FOUNDATIONS
Additional Exercises
89.
Because each pill is 4.0% Lipitor by mass, for every 100.0 g of pills, there are 4.0 g of Lipitor
present. Note that 100 pills is assumed to be an exact number.
100 pills ×
2.5 g 4.0 g Lipitor
1 kg
×
×
= 0.010 kg Lipitor
pill
100.0 g pills 1000 g
4 qt
1L
×
= 477 L
gal
1.057 qt
90.
126 gal ×
91.
100 cm  
100 cm 

9
3
Total volume =  200. m ×
 ×  300. m ×
 × 4.0 cm = 2.4 × 10 cm
m  
m 

Volume of topsoil covered by 1 bag =

10. ft 2 ×

2.4 × 109 cm3 ×
92.
2
 12 in 
 2.54 cm 

 × 

 ft 
 in 
1 bag
2.4 × 10 cm
4
3
2
2.54 cm 

4
3
 × 1.0 in ×
 = 2.4 × 10 cm
in



= 1.0 × 105 bags topsoil
a. No; if the volumes were the same, then the gold idol would have a much greater mass
because gold is much more dense than sand.
b. Mass = 1.0 L ×
1 kg
1 9.32 g
1 000 cm 3
= 19.32 kg (= 42.59 lb)
×
×
3
1000 g
L
cm
It wouldn't be easy to play catch with the idol because it would have a mass of over 40
pounds.
93.
1 light year = 1 yr ×
365 day
24 h
60 min
60 s
186,000 mi
×
×
×
×
yr
day
h
min
s
= 5.87 × 1012 miles
9.6 parsecs ×
94.
1s ×
3.26 light yr
5.87 × 1012 mi
1.609 km 1000 m
= 3.0 × 1017 m
×
×
×
parsec
light yr
mi
km
1 min
1h
65 mi 5280 ft
×
×
×
= 95.3 ft = 100 ft
60 s
60 min
h
mi
If you take your eyes off the road for one second traveling at 65 mph, your car travels
approximately 100 feet.
95.
a. 0.25 lb ×
453.6 g 1.0 g trytophan
×
= 1.1 g tryptophan
lb
100.0 g turkey
CHAPTER 1
CHEMICAL FOUNDATIONS
b. 0.25 qt ×
96.
19
0.9463 L 1.04 kg
1000 kg
2.0 g tryptophan
×
×
×
qt
L
kg
100.0 g milk
= 4.9 g tryptophan
A chemical change involves the change of one or more substances into other substances
through a reorganization of the atoms. A physical change involves the change in the form of
a substance, but not its chemical composition.
a. physical change (Just smaller pieces of the same substance.)
b. chemical change (Chemical reactions occur.)
c. chemical change (Bonds are broken.)
d. chemical change (Bonds are broken.)
e. physical change (Water is changed from a liquid to a gas.)
f.
97.
physical change (Chemical composition does not change.)
18.5 cm ×
10.0 o F
= 35.2°F increase; Tfinal = 98.6 + 35.2 = 133.8°F
5.25 cm
Tc = 5/9 (133.8 – 32) = 56.56°C
98.
Massbenzene = 58.80 g − 25.00 g = 33.80 g; Vbenzene = 33.80 g ×
Vsolid = 50.0 cm3 − 38.4 cm3 = 11.6 cm3; density =
99.
1 cm 3
= 38.4 cm3
0.880 g
25.00 g
= 2.16 g/cm3
11.6 cm 3
a. Volume × density = mass; the orange block is more dense. Because mass (orange) >
mass (blue) and because volume (orange) < volume (blue), the density of the orange
block must be greater to account for the larger mass of the orange block.
b. Which block is more dense cannot be determined. Because mass (orange) > mass (blue)
and because volume (orange) > volume (blue), the density of the orange block may or
may not be larger than the blue block. If the blue block is more dense, its density cannot
be so large that its mass is larger than the orange block’s mass.
c. The blue block is more dense. Because mass (blue) = mass (orange) and because volume
(blue) < volume (orange), the density of the blue block must be larger in order to equate
the masses.
d. The blue block is more dense. Because mass (blue) > mass (orange) and because the
volumes are equal, the density of the blue block must be larger in order to give the blue
block the larger mass.
100.
Circumference = c = 2πr; V =
4π r 3
4π  c

=
3
3  2π
3

c3
 =
6π 2

20
CHAPTER 1
Largest density =
5.25 oz
(9.00 in )
5.25 oz
=
3
12.3 in
3
CHEMICAL FOUNDATIONS
0.427 oz
=
in 3
6π 2
Smallest density =
Maximum range is:
101.
5.00 oz
(9.25 in ) 3
6π 2
dmax =
5.00 oz
13.4 in
(0.373 − 0.427) oz
in 3
V = Vfinal − Vinitial; d =
dmax =
=
9.8 cm − 6.4 cm
=
0.73 oz
in 3
or 0.40 ±0.03 oz/in3 (Uncertainty is in 2nd decimal
place.)
28.90 g
3
3
3
=
28.90 g
3.4 cm
3
= 8.5 g/cm3
mass max
; we get Vmin from 9.7 cm3 − 6.5 cm3 = 3.2 cm3.
Vmin
28.93 g
3.2 cm
3
=
9.0 g
cm
3
; dmin =
mass min
28.87 g
8.0 g
=
=
3
3
Vmax
9.9 cm − 6.3 cm
cm 3
The density is 8.5 ±0.5 g/cm3.
102.
We need to calculate the maximum and minimum values of the density, given the uncertainty
in each measurement. The maximum value is:
dmax =
19.625 g + 0.002 g
25.00 cm − 0.03 cm
3
3
=
19.627 g
24.97 cm
3
= 0.7860 g/cm3
The minimum value of the density is:
dmin =
19.625 g − 0.002 g
25.00 cm + 0.03 cm
3
3
=
19.623 g
25.03 cm 3
= 0.7840 g/cm3
The density of the liquid is between 0.7840 and 0.7860 g/cm3. These measurements are
sufficiently precise to distinguish between ethanol (d = 0.789 g/cm3) and isopropyl alcohol
(d = 0.785 g/cm3).
ChemWork Problems
The answers to the problems 103-108 (or a variation to these problems) are found in OWL. These
problems are also assignable in OWL.
Challenge Problems
109.
In a subtraction, the result gets smaller, but the uncertainties add. If the two numbers are very
close together, the uncertainty may be larger than the result. For example, let’s assume we
want to take the difference of the following two measured quantities, 999,999 ±2 and 999,996
±2. The difference is 3 ±4. Because of the uncertainty, subtracting two similar numbers is
poor practice.
CHAPTER 1
110.
CHEMICAL FOUNDATIONS
21
In general, glassware is estimated to one place past the markings.
a.
128.7 mL glassware
b.
18 mL glassware
130
c. 23.45 mL glassware
24
30
129
20
128
read to tenth’s place
23
10
127
read to one’s place
read to two decimal places
Total volume = 128.7 + 18 + 23.45 = 170.15 = 170. (Due to 18, the sum would be
known only to the ones place.)
111.
112.
a.
2.70 − 2.64
× 100 = 2%
2.70
c.
1.000 − 0.9981
0.002
× 100 =
× 100 = 0.2%
1.000
1.000
b.
| 16.12 − 16.48 |
× 100 = 2.2%
16.12
a. At some point in 1982, the composition of the metal used in minting pennies was
changed because the mass changed during this year (assuming the volume of the pennies
were constant).
b. It should be expressed as 3.08 ±0.05 g. The uncertainty in the second decimal place will
swamp any effect of the next decimal places.
113.
Heavy pennies (old): mean mass = 3.08 ±0.05 g
Light pennies (new): mean mass =
(2.467 + 2.545 + 2.518)
= 2.51 ±0.04 g
3
Because we are assuming that volume is additive, let’s calculate the volume of 100.0 g of
each type of penny, then calculate the density of the alloy. For 100.0 g of the old pennies, 95
g will be Cu (copper) and 5 g will be Zn (zinc).
1 cm 3
1 cm 3
V = 95 g Cu ×
= 11.3 cm3 (carrying one extra sig. fig.)
+ 5 g Zn ×
8.96 g
7.14 g
100. g
Density of old pennies =
= 8.8 g/cm3
11.3 cm 3
For 100.0 g of new pennies, 97.6 g will be Zn and 2.4 g will be Cu.
1 cm 3
1 cm 3
+ 97.6 g Zn ×
= 13.94 cm3 (carrying one extra sig. fig.)
8.96 g
7.14 g
100. g
Density of new pennies =
= 7.17 g/cm3
3
13.94 cm
V = 2.4 g Cu ×
22
CHAPTER 1
d=
CHEMICAL FOUNDATIONS
mass
; because the volume of both types of pennies are assumed equal, then:
volume
3
7.17 g / cm
d new
mass new
= 0.81
=
=
3
d old
massold
8.8 g / cm
The calculated average mass ratio is:
mass new
2.51 g
=
= 0.815
3.08 g
massold
To the first two decimal places, the ratios are the same. If the assumptions are correct, then
we can reasonably conclude that the difference in mass is accounted for by the difference in
alloy used.
114.
a. At 8 a.m., approximately 57 cars pass through the intersection per hour.
b. At 12 a.m. (midnight), only 1 or 2 cars pass through the intersection per hour.
c. Traffic at the intersection is limited to less than 10 cars per hour from 8 p.m. to 5 a.m.
Starting at 6 a.m., there is a steady increase in traffic through the intersection, peaking at
8 a.m. when approximately 57 cars pass per hour. Past 8 a.m. traffic moderates to about
40 cars through the intersection per hour until noon, and then decreases to 21 cars per
hour by 3 p.m. Past 3 p.m. traffic steadily increases to a peak of 52 cars per hour at 5
p.m., and then steadily decreases to the overnight level of less than 10 cars through the
intersection per hour.
d. The traffic pattern through the intersection is directly related to the work schedules of the
general population as well as to the store hours of the businesses in downtown.
e. Run the same experiment on a Sunday, when most of the general population doesn’t
work and when a significant number of downtown stores are closed in the morning.
115.
Let x = mass of copper and y = mass of silver.
105.0 g = x + y and 10.12 mL =
x
y
; solving:
+
8.96 10.5
x
105.0 − x 

+
10.12 =
 × 8.96 × 10.5, 952.1 = (10.5)x + 940.8 − (8.96)x
8.96
10.5 

(carrying 1 extra sig. fig.)
11.3 = (1.54)x, x = 7.3 g; mass % Cu =
116.
7.3 g
× 100 = 7.0% Cu
105.0 g
a.
2 compounds
compound and element (diatomic)
CHAPTER 1
CHEMICAL FOUNDATIONS
23
b.
gas element (monoatomic)
liquid element
atoms/molecules far apart;
random order; takes volume
of container
117.
atoms/molecules close
together; somewhat
ordered arrangement;
takes volume of container
solid element
atoms/molecules
close together;
ordered arrangement;
has its own volume
a. One possibility is that rope B is not attached to anything and rope A and rope C are
connected via a pair of pulleys and/or gears.
b. Try to pull rope B out of the box. Measure the distance moved by C for a given
movement of A. Hold either A or C firmly while pulling on the other rope.
118.
The bubbles of gas is air in the sand that is escaping; methanol and sand are not reacting. We
will assume that the mass of trapped air is insignificant.
Mass of dry sand = 37.3488 g − 22.8317 g = 14.5171 g
Mass of methanol = 45.2613 g − 37.3488 g = 7.9125 g
Volume of sand particles (air absent) = volume of sand and methanol − volume of methanol
Volume of sand particles (air absent) = 17.6 mL − 10.00 mL = 7.6 mL
Density of dry sand (air present) =
Density of methanol =
14.5171 g
= 1.45 g/mL
10.0 mL
7.9125 g
= 0.7913 g/mL
10.00 mL
Density of sand particles (air absent) =
14.5171 g
= 1.9 g/mL
7.6 mL
Integrative Problems
119.
2.97 × 108 persons × 0.0100 = 2.97 × 106 persons contributing
$4.75 × 108
2.97 × 10 persons
6
= $160./person;
$160.
20 nickels
×
= 3.20 × 103 nickels/person
person
$1
24
CHAPTER 1
CHEMICAL FOUNDATIONS
1£
$160.
×
= 85.6 £/person
person $1.869
120.
22610 kg
m3
×
1000 g
1 m3
= 22.61 g/cm3
×
6
3
kg
1 × 10 cm
Volume of block = 10.0 cm × 8.0 cm × 9.0 cm = 720 cm3;
121.
At 200.0°F: TC =
5
9
At −100.0°F: TC =
22.61 g
× 720 cm3 = 1.6 × 104 g
3
cm
(200.0°F − 32°F) = 93.33°C; TK = 93.33 + 273.15 = 366.48 K
5
9
(−100.0°F − 32°F) = −73.33°C; TK = −73.33°C + 273.15 = 199.82 K
∆T(°C) = [93.33°C − (−73.33°C)] = 166.66°C; ∆T(K) = (366.48 K −199.82 K) = 166.66 K
The “300 Club” name only works for the Fahrenheit scale; it does not hold true for the
Celsius and Kelvin scales.
CHAPTER 2
ATOMS, MOLECULES, AND IONS
Questions
16.
Some elements exist as molecular substances. That is, hydrogen normally exists as H2
molecules, not single hydrogen atoms. The same is true for N2, O2, F2, Cl2, etc.
17.
A compound will always contain the same numbers (and types) of atoms. A given amount of
hydrogen will react only with a specific amount of oxygen. Any excess oxygen will remain
unreacted.
18.
The halogens have a high affinity for electrons, and one important way they react is to form
anions of the type X−. The alkali metals tend to give up electrons easily and in most of their
compounds exist as M+ cations. Note: These two very reactive groups are only one electron
away (in the periodic table) from the least reactive family of elements, the noble gases.
19.
Law of conservation of mass: Mass is neither created nor destroyed. The total mass before a
chemical reaction always equals the total mass after a chemical reaction.
Law of definite proportion: A given compound always contains exactly the same proportion
of elements by mass. For example, water is always 1 g H for every 8 g oxygen.
Law of multiple proportions: When two elements form a series of compounds, the ratios of
the mass of the second element that combine with 1 g of the first element always can be
reduced to small whole numbers: For CO2 and CO discussed in Section 2.2, the mass ratios of
oxygen that react with 1 g carbon in each compound are in a 2 : 1 ratio.
20.
a. The smaller parts are electrons and the nucleus. The nucleus is broken down into protons
and neutrons, which can be broken down into quarks. For our purpose, electrons,
neutrons, and protons are the key smaller parts of an atom.
b. All atoms of hydrogen have 1 proton in the nucleus. Different isotopes of hydrogen have
0, 1, or 2 neutrons in the nucleus. Because we are talking about atoms, this implies a
neutral charge, which dictates 1 electron present for all hydrogen atoms. If charged ions
were included, then different ions/atoms of H could have different numbers of electrons.
c. Hydrogen atoms always have 1 proton in the nucleus, and helium atoms always have 2
protons in the nucleus. The number of neutrons can be the same for a hydrogen atom and
a helium atom. Tritium (3H) and 4He both have 2 neutrons. Assuming neutral atoms, then
the number of electrons will be 1 for hydrogen and 2 for helium.
d. Water (H2O) is always 1 g hydrogen for every 8 g of O present, whereas H2O2 is always 1
g hydrogen for every 16 g of O present. These are distinctly different compounds, each
with its own unique relative number and types of atoms present.
25
26
CHAPTER 2 ATOMS, MOLECULES, AND IONS
e. A chemical equation involves a reorganization of the atoms. Bonds are broken between
atoms in the reactants, and new bonds are formed in the products. The number and types
of atoms between reactants and products do not change. Because atoms are conserved in
a chemical reaction, mass is also conserved.
21.
J. J. Thomson’s study of cathode-ray tubes led him to postulate the existence of negatively
charged particles that we now call electrons. Thomson also postulated that atoms must
contain positive charge in order for the atom to be electrically neutral. Ernest Rutherford and
his alpha bombardment of metal foil experiments led him to postulate the nuclear atom−an
atom with a tiny dense center of positive charge (the nucleus) with electrons moving about
the nucleus at relatively large distances away; the distance is so large that an atom is mostly
empty space.
22.
The atom is composed of a tiny dense nucleus containing most of the mass of the atom. The
nucleus itself is composed of neutrons and protons. Neutrons have a mass slightly larger than
that of a proton and have no charge. Protons, on the other hand, have a 1+ relative charge as
compared to the 1– charged electrons; the electrons move about the nucleus at relatively large
distances. The volume of space that the electrons move about is so large, as compared to the
nucleus, that we say an atom is mostly empty space.
23.
The number and arrangement of electrons in an atom determine how the atom will react with
other atoms, i.e., the electrons determine the chemical properties of an atom. The number of
neutrons present determines the isotope identity and the mass number.
24.
Density = mass/volume; if the volumes are assumed equal, then the much more massive
proton would have a much larger density than the relatively light electron.
25.
For lighter, stable isotopes, the number of protons in the nucleus is about equal to the number
of neutrons. When the number of protons and neutrons is equal to each other, the mass
number (protons + neutrons) will be twice the atomic number (protons). Therefore, for
lighter isotopes, the ratio of the mass number to the atomic number is close to 2. For
example, consider 28Si, which has 14 protons and (28 – 14 =) 14 neutrons. Here, the mass
number to atomic number ratio is 28/14 = 2.0. For heavier isotopes, there are more neutrons
than protons in the nucleus. Therefore, the ratio of the mass number to the atomic number
increases steadily upward from 2 as the isotopes get heavier and heavier. For example, 238 U
has 92 protons and (238 – 92 =) 146 neutrons. The ratio of the mass number to the atomic
number for 238U is 238/92 = 2.6.
26.
Some properties of metals are
(1) conduct heat and electricity;
(2) malleable (can be hammered into sheets);
(3) ductile (can be pulled into wires);
(4) lustrous appearance;
(5) form cations when they form ionic compounds.
Nonmetals generally do not have these properties, and when they form ionic compounds,
nonmetals always form anions.
CHAPTER 2
ATOMS, MOLECULES, AND IONS
27
27.
Carbon is a nonmetal. Silicon and germanium are called metalloids because they exhibit both
metallic and nonmetallic properties. Tin and lead are metals. Thus metallic character
increases as one goes down a family in the periodic table. The metallic character decreases
from left to right across the periodic table.
28.
a. A molecule has no overall charge (an equal number of electrons and protons are present).
Ions, on the other hand, have extra electrons added or removed to form anions (negatively
charged ions) or cations (positively charged ions).
b. The sharing of electrons between atoms is a covalent bond. An ionic bond is the force of
attraction between two oppositely charged ions.
c. A molecule is a collection of atoms held together by covalent bonds. A compound is
composed of two or more different elements having constant composition. Covalent
and/or ionic bonds can hold the atoms together in a compound. Another difference is that
molecules do not necessarily have to be compounds. H2 is two hydrogen atoms held
together by a covalent bond. H2 is a molecule, but it is not a compound; H2 is a diatomic
element.
d. An anion is a negatively charged ion; e.g., Cl−, O2−, and SO42− are all anions. A cation is a
positively charged ion, e.g., Na+, Fe3+, and NH4+ are all cations.
29.
a. This represents ionic bonding. Ionic bonding is the electrostatic attraction between
anions and cations.
b. This represents covalent bonding where electrons are shared between two atoms. This
could be the space-filling model for H2O or SF2 or NO2, etc.
30.
Natural niacin and commercially produced niacin have the exact same formula of C6H5NO2.
Therefore, both sources produce niacin having an identical nutritional value. There may be
other compounds present in natural niacin that would increase the nutritional value, but the
nutritional value due to just niacin is identical to the commercially produced niacin.
31.
Statements a and b are true. Counting over in the periodic table, element 118 will be the next
noble gas (a nonmetal). For statement c, hydrogen has mostly nonmetallic properties. For
statement d, a family of elements is also known as a group of elements. For statement e, two
items are incorrect. When a metal reacts with a nonmetal, an ionic compound is produced,
and the formula of the compound would be AX2 (alkaline earth metals form 2+ ions and halogens form 1– ions in ionic compounds). The correct statement would be: When an alkaline
earth metal, A, reacts with a halogen, X, the formula of the ionic compound formed should be
AX2.
32.
a. Dinitrogen monoxide is correct. N and O are both nonmetals, resulting in a covalent
compound. We need to use the covalent rules of nomenclature. The other two names are
for ionic compounds.
b. Copper(I) oxide is correct. With a metal in a compound, we have an ionic compound.
Because copper, like most transition metals, forms at least a couple of different stable
charged ions in compounds, we must indicate the charge on copper in the name. Copper
oxide could be CuO or Cu2O, hence why we must give the charge of most transition
28
CHAPTER 2 ATOMS, MOLECULES, AND IONS
metal compounds. Dicopper monoxide is the name if this were a covalent compound,
which it is not.
c. Lithium oxide is correct. Lithium forms 1+ charged ions in stable ionic compounds.
Because lithium is assumed to form 1+ ions in compounds, we do not need to indicate the
charge of the metal ion in the compound. Dilithium monoxide would be the name if Li2O
were a covalent compound (a compound composed of only nonmetals).
Exercises
Development of the Atomic Theory
33.
a. The composition of a substance depends on the numbers of atoms of each element
making up the compound (depends on the formula of the compound) and not on the
composition of the mixture from which it was formed.
b. Avogadro’s hypothesis (law) implies that volume ratios are proportional to molecule
ratios at constant temperature and pressure. H2(g) + Cl2(g) → 2 HCl(g). From the
balanced equation, the volume of HCl produced will be twice the volume of H2 (or Cl2)
reacted.
34.
Avogadro’s hypothesis (law) implies that volume ratios are equal to molecule ratios at
constant temperature and pressure. Here, 1 volume of N2 reacts with 3 volumes of H2 to
produce 2 volumes of the gaseous product or in terms of molecule ratios:
1 N2 + 3 H2 → 2 product
In order for the equation to be balanced, the product must be NH3.
35.
From the law of definite proportions, a given compound always contains exactly the same
proportion of elements by mass. The first sample of chloroform has a total mass of 12.0 g C
+ 106.4 g Cl + 1.01 g H = 119.41 g (carrying extra significant figures). The mass percent of
carbon in this sample of chloroform is:
12.0 g C
× 100 = 10.05% C by mass
119.41 g total
From the law of definite proportions, the second sample of chloroform must also contain
10.05% C by mass. Let x = mass of chloroform in the second sample:
30.0 g C
× 100 = 10.05, x = 299 g chloroform
x
36.
A compound will always have a constant composition by mass. From the initial data given,
the mass ratio of H : S : O in sulfuric acid (H2SO4) is:
2.02 32.07 64.00
= 1 : 15.9 : 31.7
:
:
2.02
2.02
2.02
If we have 7.27 g H, then we will have 7.27 × 15.9 = 116 g S and 7.27 × 31.7 = 230. g O in
the second sample of H2SO4.
CHAPTER 2
37.
ATOMS, MOLECULES, AND IONS
29
Hydrazine: 1.44 × 10 −1 g H/g N; ammonia: 2.16 × 10 −1 g H/g N; hydrogen azide:
2.40 × 10 −2 g H/g N. Let's try all of the ratios:
0.216
3
0.144
0.216
0.0240
= 6.00;
= 9.00;
= 1.00;
= 1.50 =
0.0240
0.0240
0.0240
0.144
2
All the masses of hydrogen in these three compounds can be expressed as simple wholenumber ratios. The g H/g N in hydrazine, ammonia, and hydrogen azide are in the ratios
6 : 9 : 1.
38.
The law of multiple proportions does not involve looking at the ratio of the mass of one
element with the total mass of the compounds. To illustrate the law of multiple proportions,
we compare the mass of carbon that combines with 1.0 g of oxygen in each compound:
compound 1:
27.2 g C and 72.8 g O (100.0 − 27.2 = mass O)
compound 2:
42.9 g C and 57.1 g O (100.0 − 42.9 = mass O)
The mass of carbon that combines with 1.0 g of oxygen is:
compound 1:
27.2 g C
= 0.374 g C/g O
72.8 g O
compound 2:
42.9 g C
= 0.751 g C/g O
57.1 g O
0.751
2
= ; this supports the law of multiple proportions because this carbon ratio is a whole
0.374
1
number.
39.
For CO and CO2, it is easiest to concentrate on the mass of oxygen that combines with 1 g of
carbon. From the formulas (two oxygen atoms per carbon atom in CO2 versus one oxygen
atom per carbon atom in CO), CO2 will have twice the mass of oxygen that combines per
gram of carbon as compared to CO. For CO2 and C3O2, it is easiest to concentrate on the
mass of carbon that combines with 1 g of oxygen. From the formulas (three carbon atoms per
two oxygen atoms in C3O2 versus one carbon atom per two oxygen atoms in CO2), C3O2 will
have three times the mass of carbon that combines per gram of oxygen as compared to CO2.
As expected, the mass ratios are whole numbers as predicted by the law of multiple
proportions.
40.
Compound I:
14.0 g R
4.67 g R
=
;
3.00 g Q
1.00 g Q
compound II:
7.00 g R
1.56 g R
=
4.50 g Q 1.00 g Q
4.67
= 2.99 ≈ 3
1.56
As expected from the law of multiple proportions, this ratio is a small whole number.
The ratio of the masses of R that combine with 1.00 g Q is:
Because compound I contains three times the mass of R per gram of Q as compared with
compound II (RQ), the formula of compound I should be R3Q.
30
CHAPTER 2 ATOMS, MOLECULES, AND IONS
41.
Mass is conserved in a chemical reaction because atoms are conserved. Chemical reactions
involve the reorganization of atoms, so formulas change in a chemical reaction, but the
number and types of atoms do not change. Because the atoms do not change in a chemical
reaction, mass must not change. In this equation we have two oxygen atoms and four
hydrogen atoms both before and after the reaction occurs.
42.
Mass is conserved in a chemical reaction.
Mass:
ethanol + oxygen → water + carbon dioxide
46.0 g
96.0 g
54.0 g
?
Mass of reactants = 46.0 + 96.0 = 142.0 g = mass of products
142.0 g = 54.0 g + mass of CO2, mass of CO2 = 142.0 – 54.0 = 88.0 g
43.
To get the atomic mass of H to be 1.00, we divide the mass of hydrogen that reacts with 1.00
0.126
g of oxygen by 0.126; that is,
= 1.00. To get Na, Mg, and O on the same scale, we do
0.126
the same division.
Na:
2.875
1.500
1.00
= 22.8; Mg:
= 11.9; O:
= 7.94
0.126
0.126
0.126
H
O
Na
Mg
Relative value
1.00
7.94
22.8
11.9
Accepted value
1.008
16.00
22.99
24.31
For your information, the atomic masses of O and Mg are incorrect. The atomic masses of H
and Na are close to the values given in the periodic table. Something must be wrong about
the assumed formulas of the compounds. It turns out the correct formulas are H 2O, Na2O,
and MgO. The smaller discrepancies result from the error in the assumed atomic mass of H.
44.
If the formula is InO, then one atomic mass of In would combine with one atomic mass of O,
or:
A
4.784 g In
=
, A = atomic mass of In = 76.54
16.00
1.000 g O
If the formula is In2O3, then two times the atomic mass of In will combine with three times
the atomic mass of O, or:
2A
4.784 g In
=
, A = atomic mass of In = 114.8
(3)16.00
1.000 g O
The latter number is the atomic mass of In used in the modern periodic table.
The Nature of the Atom
45.
From section 2.5, the nucleus has “a diameter of about 10−13 cm” and the electrons “move
about the nucleus at an average distance of about 10−8 cm from it.” We will use these
CHAPTER 2
ATOMS, MOLECULES, AND IONS
31
statements to help determine the densities. Density of hydrogen nucleus (contains one proton
only):
Vnucleus =
4 3
4
π r = (3.14) (5 × 10 −14 cm)3 = 5 × 10 − 40 cm 3
3
3
1.67 × 10 −24 g
d = density =
5 × 10
− 40
cm
= 3 × 1015 g/cm 3
3
Density of H atom (contains one proton and one electron):
Vatom =
d=
46.
4
(3.14) (1 × 10 −8 cm) 3 = 4 × 10 − 24 cm 3
3
1.67 × 10 −24 g + 9 × 10 −28 g
4 × 10 − 24 cm 3
= 0.4 g/cm 3
Because electrons move about the nucleus at an average distance of about 1 × 10 −8 cm, the
diameter of an atom will be about 2 × 10 −8 cm. Let's set up a ratio:
diameter of nucleus
1 mm
1 × 10 −13 cm
; solving:
=
=
diameter of atom
diameter of model
2 × 10 −8 cm
diameter of model = 2 × 105 mm = 200 m
1 electron charge
47.
5.93 × 10 −18 C ×
48.
First, divide all charges by the smallest quantity, 6.40 × 10 −13 .
1.602 × 10 −19 C
2.56 × 10 −12
6.40 × 10
−13
= 4.00;
= 37 negative (electron) charges on the oil drop
7.68
3.84
= 12.0;
= 6.00
0.640
0.640
Because all charges are whole-number multiples of 6.40 × 10 −13 zirkombs, the charge on one
electron could be 6.40 × 10 −13 zirkombs. However, 6.40 × 10 −13 zirkombs could be the
charge of two electrons (or three electrons, etc.). All one can conclude is that the charge of
an electron is 6.40 × 10 −13 zirkombs or an integer fraction of 6.40 × 10 −13 zirkombs.
49.
sodium−Na; radium−Ra; iron−Fe; gold−Au; manganese−Mn; lead−Pb
50.
fluorine−F; chlorine−Cl; bromine−Br; sulfur−S; oxygen−O; phosphorus−P
51.
Sn−tin; Pt−platinum; Hg−mercury; Mg−magnesium; K−potassium; Ag−silver
52.
As−arsenic; I−iodine; Xe−xenon; He−helium; C−carbon; Si−silicon
53.
a. Metals: Mg, Ti, Au, Bi, Ge, Eu, and Am. Nonmetals: Si, B, At, Rn, and Br.
32
CHAPTER 2 ATOMS, MOLECULES, AND IONS
b. Si, Ge, B, and At. The elements at the boundary between the metals and the nonmetals
are B, Si, Ge, As, Sb, Te, Po, and At. Aluminum has mostly properties of metals, so it is
generally not classified as a metalloid.
54.
a. The noble gases are He, Ne, Ar, Kr, Xe, and Rn (helium, neon, argon, krypton, xenon,
and radon). Radon has only radioactive isotopes. In the periodic table, the whole number
enclosed in parentheses is the mass number of the longest-lived isotope of the element.
b. Promethium (Pm) has only radioactive isotopes.
55.
56.
a. transition metals
b. alkaline earth metals
d. noble gases
e. halogens
c. alkali metals
Use the periodic table to identify the elements.
a. Cl; halogen
b. Be; alkaline earth metal
c. Eu; lanthanide metal
d. Hf; transition metal
e. He; noble gas
f.
U; actinide metal
g. Cs; alkali metal
57.
58.
a.
Element 8 is oxygen. A = mass number = 9 + 8 = 17;
b.
Chlorine is element 17.
d.
Z = 26; A = 26 + 31 = 57;
f.
Lithium is element 3.
10
5B
60
27 Co
e. Iodine is element 53.
131
53 I
Li
c.
23
12 Mg
d.
132
53 I
e.
47
20 Ca
58
27 Co
f.
65
29 Cu
Z is the atomic number and is equal to the number of protons in the nucleus. A is the mass
number and is equal to the number of protons plus neutrons in the nucleus. X is the symbol
of the element. See the front cover of the text which has a listing of the symbols for the
various elements and corresponding atomic number or see the periodic table on the cover to
determine the identity of the various atoms. Because all of the atoms have equal numbers of
protons and electrons, each atom is neutral in charge.
a. 23
11 Na
60.
7
3
57
26 Fe
c. Cobalt is element 27.
a. Cobalt is element 27. A = mass number = 27 + 31 = 58;
b.
59.
37
17 Cl
17
8O
b. 199 F
c. 168 O
The atomic number for carbon is 6. 14C has 6 protons, 14 − 6 = 8 neutrons, and 6 electrons in
the neutral atom. 12C has 6 protons, 12 – 6 = 6 neutrons, and 6 electrons in the neutral atom.
The only difference between an atom of 14C and an atom of 12C is that 14C has two additional
neutrons.
CHAPTER 2
61.
62.
63.
64.
65.
a.
ATOMS, MOLECULES, AND IONS
33
79
35 Br:
35 protons, 79 – 35 = 44 neutrons. Because the charge of the atom is neutral,
the number of protons = the number of electrons = 35.
b.
81
35 Br:
c.
239
94
d.
133
55 Cs:
e.
3
1 H:
f.
56
26 Fe:
26 protons, 30 neutrons, 26 electrons
a.
235
92 U:
92 p, 143 n, 92 e
d.
208
82 Pb:
35 protons, 46 neutrons, 35 electrons
Pu: 94 protons, 145 neutrons, 94 electrons
55 protons, 78 neutrons, 55 electrons
1 proton, 2 neutrons, 1 electron
82 p, 126 n, 82 e
b.
27
13 Al:
13 p, 14 n, 13 e
c.
57
26 Fe:
26 p, 31 n, 26 e
e.
86
37 Rb:
37 p, 49 n, 37 e
f.
41
20 Ca:
20 p, 21 n, 20 e
a. Ba is element 56. Ba2+ has 56 protons, so Ba2+ must have 54 electrons in order to have a
net charge of 2+.
b.
c.
d.
e.
f.
g.
Zn is element 30. Zn2+ has 30 protons and 28 electrons.
N is element 7. N3− has 7 protons and 10 electrons.
Rb is element 37, Rb+ has 37 protons and 36 electrons.
Co is element 27. Co3+ has 27 protons and 24 electrons.
Te is element 52. Te2− has 52 protons and 54 electrons.
Br is element 35. Br− has 35 protons and 36 electrons.
a.
24
Mg: 12
12
b.
24
Mg2+:
12
12 p, 12 n, 10 e
c.
59
Co2+:
27
d.
59
Co3+:
27
27 p, 32 n, 24 e
e.
59
Co:
27
f.
79
Se:
34
34 p, 45 n, 34 e
g.
79 2−
Se :
34
34 p, 45 n, 36 e
h.
63
Ni:
28
28 p, 35 n, 28 e
i.
59 2+
Ni :
28
28 p, 31 n, 26 e
protons, 12 neutrons, 12 electrons
27 p, 32 n, 25 e
27 p, 32 n, 27 e
Atomic number = 63 (Eu); net charge = +63 − 60 = 3+; mass number = 63 + 88 = 151;
symbol:
151
3+
63 Eu
Atomic number = 50 (Sn); mass number = 50 + 68 = 118; net charge = +50 − 48 = 2+;
symbol:
118
2+
50 Sn
34
66.
CHAPTER 2 ATOMS, MOLECULES, AND IONS
Atomic number = 16 (S); net charge = +16 − 18 = 2−; mass number = 16 + 18 = 34;
symbol:
34 2−
16 S
Atomic number = 16 (S); net charge = +16 − 18 = 2−; mass number = 16 + 16 = 32;
symbol:
32 2−
16 S
67.
Number of protons in
nucleus
Number of neutrons in
nucleus
Number of
electrons
Net
charge
238
92 U
92
146
92
0
40 2+
20 Ca
20
20
18
2+
51 3+
23 V
23
28
20
3+
89
39 Y
39
50
39
0
79 −
35 Br
35
44
36
1−
31 3 −
15 P
15
16
18
3−
Symbol
68.
Number of protons in
nucleus
Number of neutrons in
nucleus
Number of
electrons
Net
charge
26
27
24
2+
59 3+
26 Fe
26
33
23
3+
210 −
85 At
85
125
86
1–
27 3+
13 Al
13
14
10
3+
128 2−
52 Te
52
76
54
2–
Symbol
53 2 +
26 Fe
CHAPTER 2
69.
70.
ATOMS, MOLECULES, AND IONS
35
In ionic compounds, metals lose electrons to form cations, and nonmetals gain electrons to
form anions. Group 1A, 2A, and 3A metals form stable 1+, 2+, and 3+ charged cations,
respectively. Group 5A, 6A, and 7A nonmetals form 3−, 2−, and 1− charged anions,
respectively.
a. Lose 2 e − to form Ra2+.
b. Lose 3 e − to form In3+.
c. Gain 3 e − to form P 3− .
d. Gain 2 e − to form Te 2− .
e. Gain 1 e − to form Br−.
f.
Lose 1 e − to form Rb+.
See Exercise 69 for a discussion of charges various elements form when in ionic compounds.
a. Element 13 is Al. Al forms 3+ charged ions in ionic compounds. Al3+
b. Se2−
c. Ba2+
d. N3−
e. Fr+
f.
Br−
Nomenclature
71.
72.
73.
a. sodium bromide
b. rubidium oxide
c. calcium sulfide
d. aluminum iodide
e. SrF2
f.
g. K3N
h. Mg3P2
a. mercury(I) oxide
b. iron(III) bromide
c. cobalt(II) sulfide
d. titanium(IV) chloride
e. Sn3N2
f.
g. HgO
h. CrS3
a. cesium fluoride
b. lithium nitride
Al2Se3
CoI3
c. silver sulfide (Silver only forms stable 1+ ions in compounds, so no Roman numerals are
needed.)
d. manganese(IV) oxide
74.
75.
76.
77.
e. titanium(IV) oxide
f.
strontium phosphide
a. ZnCl2 (Zn only forms stable +2 ions in compounds, so no Roman numerals are needed.)
b. SnF4
c. Ca3N2
e. Hg2Se
f.
a. barium sulfite
b. sodium nitrite
c. potassium permanganate
d. potassium dichromate
a. Cr(OH)3
b. Mg(CN)2
c. Pb(CO3)2
d. NH4C2H3O2
a. dinitrogen tetroxide
b. iodine trichloride
c. sulfur dioxide
d. diphosphorus pentasulfide
d. Al2S3
AgI (Ag only forms stable +1 ions in compounds.)
36
78.
79.
80.
CHAPTER 2
ATOMS, MOLECULES, AND IONS
a. B2O3
b. AsF5
c. N2O
d. SCl6
a. copper(I) iodide
b. copper(II) iodide
d. sodium carbonate
e. sodium hydrogen carbonate or sodium bicarbonate
f.
tetrasulfur tetranitride
g. selenium tetrachloride
i.
barium chromate
j.
c. cobalt(II) iodide
h. sodium hypochlorite
ammonium nitrate
a. acetic acid
b. ammonium nitrite
c. cobalt(III) sulfide
d. iodine monochloride
e. lead(II) phosphate
f.
potassium chlorate
g. sulfuric acid
h. strontium nitride
i.
aluminum sulfite
j.
k. sodium chromate
l.
hypochlorous acid
tin(IV) oxide
Note: For the compounds named as acids, we assume these are dissolved in water.
81.
In the case of sulfur, SO42− is sulfate, and SO32− is sulfite. By analogy:
SeO42−: selenate; SeO32−: selenite; TeO42−: tellurate; TeO32−: tellurite
82.
From the anion names of hypochlorite (ClO−), chlorite (ClO2−), chlorate (ClO3−), and
perchlorate (ClO4−), the oxyanion names for similar iodine ions would be hypoiodite (IO−),
iodite (IO2−), iodate (IO3−), and periodate (IO4−). The corresponding acids would be
hypoiodous acid (HIO), iodous acid (HIO2), iodic acid (HIO3), and periodic acid (HIO4).
83.
a. SF2
b. SF6
c. NaH2PO4
d. Li3N
e. Cr2(CO3)3
f.
SnF2
g. NH4C2H3O2
h. NH4HSO4
i.
Co(NO3)3
l.
NaH
j.
84.
Hg2Cl2; mercury(I) exists as Hg2 ions.
a. CrO3
b. S2Cl2
d. K2HPO4
e. AlN
f.
85.
2+
k. KClO3
NH3 (Nitrogen trihydride is the systematic name.)
(NH4)2SO3
c. NiF2
g. MnS2
h. Na2Cr2O7
i.
j.
CI4
a. Na2O
b. Na2O2
c. KCN
d. Cu(NO3)2
e. SeBr4
f.
g. PbS2
h. CuCl
HIO2
i.
GaAs (We would predict the stable ions to be Ga3+ and As 3− .)
j.
CdSe (Cadmium only forms 2+ charged ions in compounds.)
k. ZnS (Zinc only forms 2+ charged ions in compounds.)
l.
HNO2
m. P2O5
CHAPTER 2
86.
87.
88.
ATOMS, MOLECULES, AND IONS
37
a. (NH4)2HPO4
b. Hg2S
c. SiO2
d. Na2SO3
e. Al(HSO4)3
f.
NCl3
g. HBr
h. HBrO2
i.
HBrO4
j. KHS
k. CaI2
l.
CsClO4
a. nitric acid, HNO3
b. perchloric acid, HClO4
d. sulfuric acid, H2SO4
e. phosphoric acid, H3PO4
c. acetic acid, HC2H3O2
a. Iron forms 2+ and 3+ charged ions; we need to include a Roman numeral for iron.
Iron(III) chloride is correct.
b. This is a covalent compound, so use the covalent rules. Nitrogen dioxide is correct.
c. This is an ionic compound, so use the ionic rules. Calcium oxide is correct. Calcium only
forms stable 2+ ions when in ionic compounds, so no Roman numeral is needed.
d. This is an ionic compound, so use the ionic rules. Aluminum sulfide is correct.
e. This is an ionic compound, so use the ionic rules. Mg is magnesium. Magnesium acetate
is correct.
f.
Phosphide is P3−, while phosphate is PO43−. Because phosphate has a 3− charge, the
charge on iron is 3+. Iron(III) phosphate is correct.
g. This is a covalent compound, so use the covalent rules. Diphosphorus pentasulfide is
correct.
h. Because each sodium is 1+ charged, we have the O22− (peroxide) ion present. Sodium
peroxide is correct. Note that sodium oxide would be Na2O.
i.
HNO3 is nitric acid, not nitrate acid. Nitrate acid does not exist.
j.
H2S is hydrosulfuric acid or dihydrogen sulfide or just hydrogen sulfide (common name).
H2SO4 is sulfuric acid.
Additional Exercises
89.
Yes, 1.0 g H would react with 37.0 g 37Cl, and 1.0 g H would react with 35.0 g 35Cl.
No, the mass ratio of H/Cl would always be 1 g H/37 g Cl for 37Cl and 1 g H/35 g Cl for 35Cl.
As long as we had pure 37Cl or pure 35Cl, the ratios will always hold. If we have a mixture
(such as the natural abundance of chlorine), the ratio will also be constant as long as the
composition of the mixture of the two isotopes does not change.
90.
Carbon (C); hydrogen (H); oxygen (O); nitrogen (N); phosphorus (P); sulfur (S)
38
CHAPTER 2
ATOMS, MOLECULES, AND IONS
For lighter elements, stable isotopes usually have equal numbers of protons and neutrons in
the nucleus; these stable isotopes are usually the most abundant isotope for each element.
Therefore, a predicted stable isotope for each element is 12C, 2H, 16O, 14N, 30P, and 32S. These
are stable isotopes except for 30P, which is radioactive. The most stable (and most abundant)
isotope of phosphorus is 31P. There are exceptions. Also, the most abundant isotope for
hydrogen is 1H; this has just a proton in the nucleus. 2H (deuterium) is stable (not
radioactive), but 1H is also stable as well as most abundant.
2+
53
26 Fe
91.
has 26 protons, 53 – 26 = 27 neutrons, and two fewer electrons than protons (24
electrons) in order to have a net charge of 2+.
92.
a. False. Neutrons have no charge; therefore, all particles in a nucleus are not charged.
b. False. The atom is best described as having a tiny dense nucleus containing most of the
mass of the atom with the electrons moving about the nucleus at relatively large distances
away; so much so that an atom is mostly empty space.
c. False. The mass of the nucleus makes up most of the mass of the entire atom.
d. True.
e. False. The number of protons in a neutral atom must equal the number of electrons.
93.
From the Na2X formula, X has a 2− charge. Because 36 electrons are present, X has 34
protons and 79 − 34 = 45 neutrons, and is selenium.
a. True. Nonmetals bond together using covalent bonds and are called covalent compounds.
b. False. The isotope has 34 protons.
c. False. The isotope has 45 neutrons.
d. False. The identity is selenium, Se.
94.
a. Fe2+: 26 protons (Fe is element 26.); protons − electrons = net charge, 26 − 2 = 24
electrons; FeO is the formula since the oxide ion has a 2− charge, and the name is
iron(II) oxide.
b. Fe3+: 26 protons; 23 electrons; Fe2O3; iron(III) oxide
c. Ba2+: 56 protons; 54 electrons; BaO; barium oxide
d. Cs+: 55 protons; 54 electrons; Cs2O; cesium oxide
e. S2−: 16 protons; 18 electrons; Al2S 3; aluminum sulfide
f.
P3−: 15 protons; 18 electrons; AlP; aluminum phosphide
g. Br−: 35 protons; 36 electrons; AlBr3; aluminum bromide
h. N3−: 7 protons; 10 electrons; AlN; aluminum nitride
CHAPTER 2
95.
ATOMS, MOLECULES, AND IONS
39
a. Pb(C2H3O2)2: lead(II) acetate
b. CuSO4: copper(II) sulfate
c. CaO: calcium oxide
d. MgSO4: magnesium sulfate
e. Mg(OH)2: magnesium hydroxide
f.
CaSO4: calcium sulfate
g. N2O: dinitrogen monoxide or nitrous oxide (common name)
96.
a. This is element 52, tellurium. Te forms stable 2! charged ions in ionic compounds (like
other oxygen family members).
b.
Rubidium. Rb, element 37, forms stable 1+ charged ions.
c.
Argon. Ar is element 18.
d.
Astatine. At is element 85.
97.
From the XBr2 formula, the charge on element X is 2+. Therefore, the element has 88
protons, which identifies it as radium, Ra. 230 − 88 = 142 neutrons.
98.
Because this is a relatively small number of neutrons, the number of protons will be very
close to the number of neutrons present. The heavier elements have significantly more
neutrons than protons in their nuclei. Because this element forms anions, it is a nonmetal and
will be a halogen because halogens form stable 1− charged ions in ionic compounds. From
the halogens listed, chlorine, with an average atomic mass of 35.45, fits the data. The two
isotopes are 35Cl and 37Cl, and the number of electrons in the 1− ion is 18. Note that because
the atomic mass of chlorine listed in the periodic table is closer to 35 than 37, we can assume
that 35Cl is the more abundant isotope. This is discussed in Chapter 3.
99.
a. Ca2+ and N3−: Ca3N2, calcium nitride
b. K+ and O2−: K2O, potassium oxide
c. Rb+ and F−: RbF, rubidium fluoride
d. Mg2+ and S2−: MgS, magnesium sulfide
e. Ba2+ and I−: BaI2, barium iodide
f.
Al3+ and Se2−: Al2Se3, aluminum selenide
g. Cs+ and P3−: Cs3P, cesium phosphide
h. In3+ and Br−: InBr3, indium(III) bromide. In also forms In+ ions, but one would predict
In3+ ions from its position in the periodic table.
100.
These compounds are similar to phosphate (PO43-− ) compounds. Na3AsO4 contains Na+ ions
and AsO43− ions. The name would be sodium arsenate. H3AsO4 is analogous to phosphoric
acid, H3PO4. H3AsO4 would be arsenic acid. Mg3(SbO4)2 contains Mg2+ ions and SbO43−
ions, and the name would be magnesium antimonate.
101.
a. Element 15 is phosphorus, P. This atom has 15 protons and 31 − 15 = 16 neutrons.
b. Element 53 is iodine, I. 53 protons; 74 neutrons
40
CHAPTER 2
ATOMS, MOLECULES, AND IONS
c. Element 19 is potassium, K. 19 protons; 20 neutrons
d. Element 70 is ytterbium, Yb. 70 protons; 103 neutrons
102.
Mass is conserved in a chemical reaction.
Mass:
chromium(III) oxide + aluminum → chromium + aluminum oxide
34.0 g
12.1 g
23.3 g
?
Mass aluminum oxide produced = (34.0 + 12.1) − 23.3 = 22.8 g
ChemWork Problems
The answers to the problems 103-108 (or a variation to these problems) are found in OWL. These
problems are also assignable in OWL.
Challenge Problems
109.
Copper (Cu), silver (Ag), and gold (Au) make up the coinage metals.
110.
Because the gases are at the same temperature and pressure, the volumes are directly
proportional to the number of molecules present. Let’s assume hydrogen and oxygen to be
monatomic gases and that water has the simplest possible formula (HO). We have the
equation:
H + O → HO
But the volume ratios are also equal to the molecule ratios, which correspond to the
coefficients in the equation:
2 H + O → 2 HO
Because atoms cannot be created nor destroyed in a chemical reaction, this is not possible. To
correct this, we can make oxygen a diatomic molecule:
2 H + O2 → 2 HO
This does not require hydrogen to be diatomic. Of course, if we know water has the formula
H2O, we get:
2 H + O2 → 2 H2O
The only way to balance this is to make hydrogen diatomic:
2 H2 + O2 → 2 H2O
111.
Avogadro proposed that equal volumes of gases (at constant temperature and pressure)
contain equal numbers of molecules. In terms of balanced equations, Avogadro’s hypothesis
(law) implies that volume ratios will be identical to molecule ratios. Assuming one molecule
of octane reacting, then 1 molecule of CxHy produces 8 molecules of CO2 and 9 molecules of
H2O. CxHy + n O2 → 8 CO2 + 9 H2O. Because all the carbon in octane ends up as carbon in
CHAPTER 2
ATOMS, MOLECULES, AND IONS
41
CO2, octane must contain 8 atoms of C. Similarly, all hydrogen in octane ends up as
hydrogen in H2O, so one molecule of octane must contain 9 × 2 = 18 atoms of H. Octane
formula = C8H18, and the ratio of C : H = 8 : 18 or 4 : 9.
112.
From Section 2.5 of the text, the average diameter of the nucleus is about 10−13 cm, and the
electrons move about the nucleus at an average distance of about 10 −8 cm . From this, the
diameter of an atom is about 2 × 10 −8 cm .
2 × 10 −8 cm
1 × 10
−13
cm
= 2 × 105;
1 mi
5280 ft
63,360 in
=
=
1 grape
1 grape
1 grape
Because the grape needs to be 2 × 105 times smaller than a mile, the diameter of the grape
would need to be 63,360/(2 × 105) ≈ 0.3 in. This is a reasonable size for a small grape.
113.
The alchemists were incorrect. The solid residue must have come from the flask.
114.
The equation for the reaction would be 2 Na(s) + Cl2(g) → 2 NaCl(s). The sodium reactant
exists as singular sodium atoms packed together very tightly and in a very organized fashion.
This type of packing of atoms represents the solid phase. The chlorine reactant exists as Cl2
molecules. In the picture of chlorine, there is a lot of empty space present. This only occurs
in the gaseous phase. When sodium and chlorine react, the ionic compound NaCl forms.
NaCl exists as separate Na+ and Cl− ions. Because the ions are packed very closely together
and are packed in a very organized fashion, NaCl is depicted in the solid phase.
115.
a. Both compounds have C2H6O as the formula. Because they have the same formula, their
mass percent composition will be identical. However, these are different compounds
with different properties because the atoms are bonded together differently. These
compounds are called isomers of each other.
b. When wood burns, most of the solid material in wood is converted to gases, which
escape. The gases produced are most likely CO2 and H2O.
c. The atom is not an indivisible particle but is instead composed of other smaller particles,
called electrons, neutrons, and protons.
d. The two hydride samples contain different isotopes of either hydrogen and/or lithium.
Although the compounds are composed of different isotopes, their properties are similar
because different isotopes of the same element have similar properties (except, of course,
their mass).
116.
Let Xa be the formula for the atom/molecule X, Yb be the formula for the atom/molecule Y,
XcYd be the formula of compound I between X and Y, and XeYf be the formula of compound
II between X and Y. Using the volume data, the following would be the balanced equations
for the production of the two compounds.
Xa + 2 Yb → 2 XcYd; 2 Xa + Yb → 2 XeYf
From the balanced equations, a = 2c = e and b = d = 2f.
Substituting into the balanced equations:
42
CHAPTER 2
ATOMS, MOLECULES, AND IONS
X2c + 2 Y2f → 2 XcY2f ; 2 X2c + Y2f → 2 X2cYf
For simplest formulas, assume that c = f = 1. Thus:
X2 + 2 Y2 → 2 XY2 and 2 X2 + Y2 → 2 X2Y
1.00
= 0.3043, y = 1.14.
1.00 + 2 y
2.00
Compound II = X2Y: If X has relative mass of 1.00,
= 0.6364, y = 1.14.
2.00 + y
Compound I = XY2: If X has relative mass of 1.00,
The relative mass of Y is 1.14 times that of X. Thus, if X has an atomic mass of 100, then Y
will have an atomic mass of 114.
117.
Most of the mass of the atom is due to the protons and the neutrons in the nucleus, and
protons and neutrons have about the same mass (1.67 × 10−24 g). The ratio of the mass of the
molecule to the mass of a nuclear particle will give a good approximation of the number of
nuclear particles (protons and neutrons) present.
7.31 × 10 −23 g
= 43.8 ≈ 44 nuclear particles
1.67 × 10 −24 g
Thus there are 44 protons and neutrons present. If the number of protons equals the number
of neutrons, we have 22 protons in the molecule. One possibility would be the molecule CO2
[6 + 2(8) = 22 protons].
118.
For each experiment, divide the larger number by the smaller. In doing so, we get:
experiment 1
experiment 2
experiment 3
X = 1.0
Y = 1.4
X = 1.0
Y = 10.5
Z = 1.0
Y = 3.5
Our assumption about formulas dictates the rest of the solution. For example, if we assume
that the formula of the compound in experiment 1 is XY and that of experiment 2 is YZ, we
get relative masses of:
X = 2.0; Y = 21; Z = 15 (= 21/1.4)
and a formula of X3Y for experiment 3 [three times as much X must be present in experiment
3 as compared to experiment 1 (10.5/3.5 = 3)].
However, if we assume the formula for experiment 2 is YZ and that of experiment 3 is XZ,
then we get:
X = 2.0; Y = 7.0; Z = 5.0 (= 7.0/1.4)
and a formula of XY3 for experiment 1. Any answer that is consistent with your initial
assumptions is correct.
CHAPTER 2
ATOMS, MOLECULES, AND IONS
43
The answer to part d depends on which (if any) of experiments 1 and 3 have a formula of XY
in the compound. If the compound in experiment 1 has a formula of XY, then:
21 g XY ×
4.2 g Y
= 19.2 g Y (and 1.8 g X)
(4.2 + 0.4) g XY
If the compound in experiment 3 has the XY formula, then:
21 g XY Η
7.0 g Y
= 16.3 g Y (and 4.7 g X)
(7.0 + 2.0) g XY
Note that it could be that neither experiment 1 nor experiment 3 has XY as the formula.
Therefore, there is no way of knowing an absolute answer here.
Integrated Problems
119.
The systematic name of Ta2O5 is tantalum(V) oxide. Tantalum is a transition metal and
requires a Roman numeral. Sulfur is in the same group as oxygen, and its most common ion
is S2–. There-fore, the formula of the sulfur analogue would be Ta2S5.
Total number of protons in Ta2O5:
Ta, Z = 73, so 73 protons × 2 = 146 protons; O, Z = 8, so 8 protons × 5 = 40 protons
Total protons = 186 protons
Total number of protons in Ta2S5:
Ta, Z = 73, so 73 protons × 2 = 146 protons; S, Z = 16, so 16 protons × 5 = 80 protons
Total protons = 226 protons
Proton difference between Ta2S5 and Ta2O5: 226 protons – 186 protons = 40 protons
120.
The cation has 51 protons and 48 electrons. The number of protons corresponds to the atomic
number. Thus this is element 51, antimony. There are 3 fewer electrons than protons.
Therefore, the charge on the cation is 3+. The anion has one-third the number of protons of
the cation, which corresponds to 17 protons; this is element 17, chlorine. The number of
electrons in this anion of chlorine is 17 + 1 = 18 electrons. The anion must have a charge of
1−.
The formula of the compound formed between Sb3+ and Cl– is SbCl3. The name of the
compound is antimony(III) chloride. The Roman numeral is used to indicate the charge on Sb
because the predicted charge is not obvious from the periodic table.
121.
Number of electrons in the unknown ion:
2.55 × 10 −26 g ×
1 kg
1 electron
= 28 electrons
×
1000 g 9.11 × 10 −31 kg
Number of protons in the unknown ion:
44
CHAPTER 2
5.34 × 10 −23 g ×
ATOMS, MOLECULES, AND IONS
1 kg
1 proton
= 32 protons
×
1000 g 1.67 × 10 − 27 kg
Therefore, this ion has 32 protons and 28 electrons. This is element number 32, germanium
(Ge). The net charge is 4+ because four electrons have been lost from a neutral germanium
atom.
The number of electrons in the unknown atom:
3.92 × 10 −26 g ×
1 kg
1 electron
= 43 electrons
×
1000 g 9.11 × 0 −31 kg
In a neutral atom, the number of protons and electrons is the same. Therefore, this is element
43, technetium (Tc).
The number of neutrons in the technetium atom:
1 proton
1 kg
9.35 × 10 −23 g ×
= 56 neutrons
×
1000 g 1.67 × 10 − 27 kg
The mass number is the sum of the protons and neutrons. In this atom, the mass number is 43
protons + 56 neutrons = 99. Thus this atom and its mass number is 99Tc.
Marathon Problem
122.
a.
For each set of data, divide the larger number by the smaller number to determine
relative masses.
0.602
= 2.04; A = 2.04 when B = 1.00
0.295
0.401
= 2.33; C = 2.33 when B = 1.00
0.172
0.374
= 1.17; C = 1.17 when A = 1.00
0.320
To have whole numbers, multiply the results by 3.
Data set 1: A = 6.1 and B = 3.0
Data set 2: C = 7.0 and B = 3.0
Data set 3: C = 3.5 and A = 3.0 or C = 7.0 and A = 6.0
Assuming 6.0 for the relative mass of A, the relative masses would be A = 6.0, B = 3.0,
and C = 7.0 (if simplest formulas are assumed).
CHAPTER 2
ATOMS, MOLECULES, AND IONS
45
b. Gas volumes are proportional to the number of molecules present. There are many
possible correct answers for the balanced equations. One such solution that fits the gas
volume data is:
6 A2 + B4
→ 4 A3B
B4 + 4 C3 → 4 BC3
3 A2 + 2 C3 → 6 AC
In any correct set of reactions, the calculated mass data must match the mass data given
initially in the problem. Here, the new table of relative masses would be:
6 (mass A 2 )
0.602
=
; mass A2 = 0.340(mass B4)
mass B 4
0.295
4 (mass C 3 )
0.401
=
; mass C3 = 0.583(mass B4)
mass B 4
0.172
2 (mass C 3 )
0.374
=
; mass A2 = 0.570(mass C3)
3 (mass A 2 )
0.320
Assume some relative mass number for any of the masses. We will assume that mass B =
3.0, so mass B4 = 4(3.0) = 12.
Mass C3 = 0.583(12) = 7.0, mass C = 7.0/3
Mass A2 = 0.570(7.0) = 4.0, mass A = 4.0/2 = 2.0
When we assume a relative mass for B = 3.0, then A = 2.0 and C = 7.0/3. The relative
masses having all whole numbers would be A = 6.0, B = 9.0, and C = 7.0.
Note that any set of balanced reactions that confirms the initial mass data is correct. This
is just one possibility.
CHAPTER 3
STOICHIOMETRY
Questions
23.
Isotope
12
C
13
C
Mass
12.0000 u
13.034 u
Abundance
98.89%
1.11%
Average mass = 0.9889 (12.0000) + 0.0111(13.034) = 12.01 u
Note: u is an abbreviation for amu (atomic mass units).
From the relative abundances, there would be 9889 atoms of 12C and 111 atoms of 13C in the
10,000 atom sample. The average mass of carbon is independent of the sample size; it will
always be 12.01 u.
Total mass = 10,000 atoms ×
12.01 u
= 1.201 × 105 u
atom
For 1 mole of carbon (6.0221 × 1023 atoms C), the average mass would still be 12.01 u.
The number of 12C atoms would be 0.9889(6.0221 × 1023) = 5.955 × 1023 atoms 12C, and the
number of 13C atoms would be 0.0111(6.0221 × 1023) = 6.68 × 1021 atoms 13C.
Total mass = 6.0221 × 1023 atoms ×
12.01 u
= 7.233 × 1024 u
atom
Total mass in g = 6.0221 × 1023 atoms ×
12.01 u
1g
×
= 12.01 g/mol
atom
6.0221 × 10 23 u
By using the carbon-12 standard to define the relative masses of all of the isotopes, as well as
to define the number of things in a mole, then each element’s average atomic mass in units of
grams is the mass of a mole of that element as it is found in nature.
24.
Consider a sample of glucose, C6H12O6. The molar mass of glucose is 180.16 g/mol. The
chemical formula allows one to convert from molecules of glucose to atoms of carbon,
hydrogen, or oxygen present and vice versa. The chemical formula also gives the mole
relationship in the formula. One mole of glucose contains 6 mol C, 12 mol H, and 6 mol O.
Thus mole conversions between molecules and atoms are possible using the chemical formula. The molar mass allows one to convert between mass and moles of compound, and
Avogadro’s number (6.022 × 1023) allows one to convert between moles of compound and
number of molecules.
46
CHAPTER 3
25.
STOICHIOMETRY
47
Avogadro’s number of dollars = 6.022 × 1023 dollars/mol dollars
1 mol dollars ×
6.022 × 10 23 dollars
mol dollars
7 × 10 9 people
= 8.6 × 1013 = 9 × 1013 dollars/person
1 trillion = 1,000,000,000,000 = 1 × 1012; each person would have 90 trillion dollars.
26.
Molar mass of CO2 = 12.01 + 2(16.00) = 44.01 g/mol
One mol of CO2 contains 6.022 × 1023 molecules of CO2, 6.022 × 1023 atoms of C, and 1.204
× 1024 atoms of O. We could also break down 1 mol of CO2 into the number of protons and
the number of electrons present (1.325 × 1025 protons and 1.325 × 1025 electrons). In order to
determine the number of neutrons present, we would need to know the isotope abundances
for carbon and oxygen.
The mass of 1 mol of CO2 would be 44.01 g. From the molar mass, one mol of CO2 would
contain 12.01 g C and 32.00 g O. We could also break down 1 mol of CO2 into the mass of
protons and mass of electrons present (22.16 g protons and 1.207 × 10−2 g electrons). This
assumes no mass loss when the individual particles come together to form the atom. This is
not a great assumption as will be discussed in Chapter 19 on Nuclear Chemistry.
27.
Only in b are the empirical formulas the same for both compounds illustrated. In b, general
formulas of X2Y4 and XY2 are illustrated, and both have XY2 for an empirical formula.
For a, general formulas of X2Y and X2Y2 are illustrated. The empirical formulas for these
two compounds are the same as the molecular formulas. For c, general formulas of XY and
XY2 are illustrated; these general formulas are also the empirical formulas. For d, general
formulas of XY4 and X2Y6 are illustrated. XY4 is also the molecular formula, but X2Y6 has
the empirical formula of XY3.
28.
The molar mass is the mass of 1 mole of the compound. The empirical mass is the mass of 1
mole of the empirical formula. The molar mass is a whole-number multiple of the empirical
mass. The masses are the same when the molecular formula = empirical formula, and the
masses are different when the two formulas are different. When different, the empirical mass
must be multiplied by the same whole number used to convert the empirical formula to the
molecular formula. For example, C6H12O6 is the molecular formula for glucose, and CH2O is
the empirical formula. The whole-number multiplier is 6. This same factor of 6 is the multiplier used to equate the empirical mass (30 g/mol) of glucose to the molar mass (180 g/mol).
29.
The mass percent of a compound is a constant no matter what amount of substance is present.
Compounds always have constant composition.
30.
A balanced equation starts with the correct formulas of the reactants and products. The coefficients necessary to balance the equation give molecule relationships as well as mole
relationships between reactants and products. The state (phase) of the reactants and products
is also given. Finally, special reaction conditions are sometimes listed above or below the
arrow. These can include special catalysts used and/or special temperatures required for a
reaction to occur.
48
CHAPTER 3
STOICHIOMETRY
31.
Only one product is formed in this representation. This product has two Y atoms bonded to
an X. The other substance present in the product mixture is just the excess of one of the
reactants (Y). The best equation has smallest whole numbers. Here, answer c would be this
smallest whole number equation (X + 2 Y → XY2). Answers a and b have incorrect
products listed, and for answer d, an equation only includes the reactants that go to produce
the product; excess reactants are not shown in an equation.
32.
A balanced equation must have the same number and types of atoms on both sides of the
equation, but it also needs to have correct formulas. The illustration has the equation as:
H + O → H2O
Under normal conditions, hydrogen gas and oxygen gas exist as diatomic molecules. So the
first change to make is to change H + O to H2 + O2. To balance this equation, we need one
more oxygen atom on the product side. Trial and error eventually gives the correct balanced
equation of:
2 H2 + O2 → 2 H2O
This equation uses the smallest whole numbers and has the same number of oxygen atoms
and hydrogen atoms on both sides of the equation (4 H + 2 O atoms). So in your drawing,
there should be two H2 molecules, 1 O2 molecule, and 2 H2O molecules.
33.
The theoretical yield is the stoichiometric amount of product that should form if the limiting
reactant is completely consumed and the reaction has 100% yield.
34.
A reactant is present in excess if there is more of that reactant present than is needed to
combine with the limiting reactant for the process. By definition, the limiting reactant cannot
be present in excess. An excess of any reactant does not affect the theoretical yield for a
process; the theoretical yield is determined by the limiting reactant.
35.
The specific information needed is mostly the coefficients in the balanced equation and the
molar masses of the reactants and products. For percent yield, we would need the actual yield
of the reaction and the amounts of reactants used.
a. Mass of CB produced = 1.00 × 104 molecules A2B2
molar mass of CB
1 mol A 2 B 2
2 mol CB
×
×
×
23
mol CB
6.022 × 10 molecules A 2 B 2 1 mol A 2 B 2
b. Atoms of A produced = 1.00 × 104 molecules A2B2 ×
c. Moles of C reacted = 1.00 × 104 molecules A2B2 ×
2 atoms A
1 molecule A 2 B 2
1 mol A 2 B 2
6.022 × 10 23 molecules A 2 B 2
2 mol C
×
1 mol A 2 B 2
actual mass
× 100; the theoretical mass of CB produced was
theoretical mass
calculated in part a. If the actual mass of CB produced is given, then the percent yield can
be determined for the reaction using the percent yield equation.
d. Percent yield =
CHAPTER 3
36.
STOICHIOMETRY
49
One method is to assume each quantity of reactant is limiting, then calculate the amount of
product that could be produced from each reactant. This gives two possible answers
(assuming two reactants). The correct answer (the amount of product that could be produced)
is always the smaller number. Even though there is enough of the other reactant to form more
product, once the smaller quantity is reached, the limiting reactant runs out, and the reaction
cannot continue.
A second method would be to pick one of the reactants and then calculate how much of the
other reactant would be required to react with all of it. How the answer compares to the
actual amount of that reactant present allows one to deduce the identity of the limiting
reactant. Once the identity is known, one would take the limiting reactant and convert it to
mass of product formed.
Exercises
Atomic Masses and the Mass Spectrometer
37.
Let A = average atomic mass
A = 0.0140(203.973) + 0.2410(205.9745) + 0.2210(206.9759) + 0.5240(207.9766)
A = 2.86 + 49.64 + 45.74 + 109.0 = 207.2 u; from the periodic table, the element is Pb.
Note: u is an abbreviation for amu (atomic mass units).
38.
Average atomic mass = A = 0.0800(45.952632) + 0.0730(46.951764) + 0.7380(47.947947)
+ 0.0550(48.947841) + 0.0540(49.944792) = 47.88 amu
This is element Ti (titanium).
39.
Let A = mass of 185Re:
186.207 = 0.6260(186.956) + 0.3740(A), 186.207 − 117.0 = 0.3740(A)
A=
40.
69.2
= 185 u (A = 184.95 u without rounding to proper significant figures.)
0.3740
Abundance 28Si = 100.00 − (4.70 + 3.09) = 92.21%; from the periodic table, the average
atomic mass of Si is 28.09 u.
28.09 = 0.9221(27.98) + 0.0470(atomic mass 29Si) + 0.0309(29.97)
Atomic mass 29Si = 29.01 u
The mass of 29Si is actually a little less than 29 u. There are other isotopes of silicon that are
considered when determining the 28.09 u average atomic mass of Si listed in the atomic table.
41.
Let x = % of 151Eu and y = % of 153Eu, then x + y = 100 and y = 100 − x.
50
CHAPTER 3
151.96 =
STOICHIOMETRY
x(150.9196) + (100 − x)(152.9209)
100
15196 = (150.9196)x + 15292.09 − (152.9209)x, −96 = −(2.0013)x
x = 48%; 48% 151Eu and 100 − 48 = 52% 153Eu
42.
If silver is 51.82% 107Ag, then the remainder is 109Ag (48.18%). Determining the atomic
mass (A) of 109Ag:
107.868 =
51.82(106.905) + 48.18(A)
100
10786.8 = 5540. + (48.18)A, A = 108.9 u = atomic mass of 109Ag
43.
There are three peaks in the mass spectrum, each 2 mass units apart. This is consistent with
two isotopes differing in mass by two mass units. The peak at 157.84 corresponds to a Br2
molecule composed of two atoms of the lighter isotope. This isotope has mass equal to
157.84/2 or 78.92. This corresponds to 79Br. The second isotope is 81Br with mass equal to
161.84/2 = 80.92. The peaks in the mass spectrum correspond to 79Br2, 79Br81Br, and 81Br2 in
order of increasing mass. The intensities of the highest and lowest masses tell us the two
isotopes are present in about equal abundance. The actual abundance is 50.68% 79Br and
49.32% 81Br.
44.
Because we are not given the relative masses of the isotopes, we need to estimate the masses
of the isotopes. A good estimate is to assume that only the protons and neutrons contribute to
the overall mass of the atom and that the atomic mass of a proton and neutron are each 1.00 u.
So the masses are about: 54Fe, 54.00 u; 56Fe, 56.00 u; 57Fe, 57.00 u; 58Fe, 58.00 u. Using
these masses, the calculated average atomic mass would be:
0.0585(54.00) + 0.9175(56.00) + 0.0212(57.00) + 0.0028(58.00) = 55.91 u
The average atomic mass listed in the periodic table is 55.85 u.
Moles and Molar Masses
45.
When more than one conversion factor is necessary to determine the answer, we will usually
put all the conversion factors into one calculation instead of determining intermediate
answers. This method reduces round-off error and is a time saver.
500. atoms Fe ×
46.
500.0 g Fe ×
1 mol Fe
6.022 × 10
23
atoms Fe
×
55.85 g Fe
= 4.64 × 10 −20 g Fe
mol Fe
1 mol Fe
= 8.953 mol Fe
55.85 g Fe
8.953 mol Fe ×
6.022 × 10 23 atoms Fe
= 5.391 × 1024 atoms Fe
mol Fe
CHAPTER 3
STOICHIOMETRY
0.200 g C
1 mol C
6.022 × 10 23 atoms C
= 1.00 × 1022 atoms C
×
×
carat
12.01 g C
mol C
47.
1.00 carat ×
48.
5.0 × 1021 atoms C ×
8.3 × 10 −3 mol C ×
49.
51
1 mol C
6.022 × 10 23 atoms C
= 8.3 × 10 −3 mol C
12.01 g C
= 0.10 g C
mol C
Al2O3: 2(26.98) + 3(16.00) = 101.96 g/mol
Na3AlF6: 3(22.99) + 1(26.98) + 6(19.00) = 209.95 g/mol
50.
HFC−134a, CH2FCF3: 2(12.01) + 2(1.008) + 4(19.00) = 102.04 g/mol
HCFC−124, CHClFCF3: 2(12.01) + 1(1.008) + 1(35.45) + 4(19.00) = 136.48 g/mol
51.
a. The formula is NH3. 14.01 g/mol + 3(1.008 g/mol) = 17.03 g/mol
b. The formula is N2H4. 2(14.01) + 4(1.008) = 32.05 g/mol
c. (NH4)2Cr2O7: 2(14.01) + 8(1.008) + 2(52.00) + 7(16.00) = 252.08 g/mol
52.
a. The formula is P4O6. 4(30.97 g/mol) + 6(16.00 g/mol) = 219.88 g/mol
b. Ca3(PO4)2: 3(40.08) + 2(30.97) + 8(16.00) = 310.18 g/mol
c. Na2HPO4: 2(22.99) + 1(1.008) + 1(30.97) + 4(16.00) = 141.96 g/mol
53.
a. 1.00 g NH3 ×
b. 1.00 g N2H4 ×
1 mol NH 3
= 0.0587 mol NH3
17.03 g NH 3
1 mol N 2 H 4
= 0.0312 mol N2H4
32.05 g N 2 H 4
c. 1.00 g (NH4)2Cr2O7 ×
54.
a. 1.00 g P4O6 ×
1 mol ( NH 4 ) 2 Cr2 O 7
= 3.97 × 10 −3 mol (NH4)2Cr2O7
252.08 g ( NH 4 ) 2 Cr2 O 7
1 mol P 4 O 6
= 4.55 × 10 −3 mol P4O6
219.88 g
b. 1.00 g Ca3(PO4)2 ×
1 mol Ca 3 (PO 4 ) 2
= 3.22 × 10 −3 mol Ca3(PO4)2
310.18 g
c. 1.00 g Na2HPO4 ×
1 mol Na 2 HPO 4
= 7.04 × 10 −3 mol Na2HPO4
141.96 g
52
55.
CHAPTER 3
a. 5.00 mol NH3 ×
b. 5.00 mol N2H4 ×
17.03 g NH 3
= 85.2 g NH3
mol NH 3
32.05 g N 2 H 4
= 160. g N2H4
mol N 2 H 4
c. 5.00 mol (NH4)2Cr2O7 ×
56.
a. 5.00 mol P4O6 ×
c. 5.00 mol Na2HPO4 ×
310.18 g
= 1.55 × 103 g Ca3(PO4)2
mol Ca 3 (PO 4 ) 2
141.96 g
= 7.10 × 102 g Na2HPO4
mol Na 2 HPO 4
Chemical formulas give atom ratios as well as mole ratios.
a. 5.00 mol NH3 ×
b. 5.00 mol N2H4 ×
1 mol N
14.01 g N
×
= 70.1 g N
mol NH 3
mol N
2 mol N
14.01 g N
×
= 140. g N
mol N 2 H 4
mol N
c. 5.00 mol (NH4)2Cr2O7 ×
58.
59.
252.08 g ( NH 4 ) 2 Cr2 O 7
= 1260 g (NH4)2Cr2O7
1 mol ( NH 4 ) 2 Cr2 O 7
219.88 g
= 1.10 × 103 g P4O6
1 mol P 4 O 6
b. 5.00 mol Ca3(PO4)2 ×
57.
STOICHIOMETRY
a. 5.00 mol P4O6 ×
2 mol N
14.01 g N
×
= 140. g N
mol ( NH 4 ) 2 Cr2 O 7
mol N
4 mol P
30.97 g P
×
= 619 g P
mol P4 O 6
mol P
b. 5.00 mol Ca3(PO4)2 ×
2 mol P
30.97 g P
×
= 310. g P
mol Ca 3 (PO 4 ) 2
mol P
c. 5.00 mol Na2HPO4 ×
1 mol P
30.97 g P
×
= 155 g P
mol Na 2 HPO 4
mol P
a. 1.00 g NH3 ×
b. 1.00 g N2H4 ×
1 mol NH 3
6.022 × 10 23 molecules NH 3
×
17.03 g NH 3
mol NH 3
= 3.54 × 1022 molecules NH3
1 mol N 2 H 4
6.022 × 10 23 molecules N 2 H 4
×
32.05 g N 2 H 4
mol N 2 H 4
= 1.88 × 1022 molecules N2H4
CHAPTER 3
STOICHIOMETRY
c. 1.00 g (NH4)2Cr2O7 ×
×
60.
53
1 mol ( NH 4 ) 2 Cr2 O 7
252.08 g ( NH 4 ) 2 Cr2 O 7
6.022 × 10 23 formula units ( NH 4 ) 2 Cr2 O 7
= 2.39 × 1021 formula units (NH4)2Cr2O7
mol ( NH 4 ) 2 Cr2 O 7
a. 1.00 g P4O6 ×
1 mol P4 O 6
6.022 × 10 23 molecules
= 2.74 × 1021 molecules P4O6
×
219.88 g
mol P4 O 6
b. 1.00 g Ca3(PO4)2 ×
1 mol Ca 3 (PO 4 ) 2
6.022 × 10 23 formula units
×
310.18 g
mol Ca 3 (PO 4 ) 2
= 1.94 × 1021 formula units Ca3(PO4)2
c. 1.00 g Na2HPO4 ×
1 mol Na 2 HPO 4
6.022 × 10 23 formula units
×
141.96 g
mol Na 2 HPO 4
= 4.24 × 1021 formula units Na2HPO4
61.
Using answers from Exercise 59:
a. 3.54 × 1022 molecules NH3 ×
b. 1.88 × 1022 molecules N2H4 ×
1 atom N
= 3.54 × 1022 atoms N
molecule NH 3
2 atoms N
= 3.76 × 1022 atoms N
molecule N 2 H 4
c. 2.39 × 1021 formula units (NH4)2Cr2O7 ×
2 atoms N
formula unit ( NH 4 ) 2 Cr2 O 7
= 4.78 × 1021 atoms N
62.
Using answers from Exercise 60:
a. 2.74 × 1021 molecules P4O6 ×
63.
4 atoms P
= 1.10 × 1022 atoms P
molecule P4 O 6
b. 1.94 × 1021 formula units Ca3(PO4)2 ×
2 atoms P
= 3.88 × 1021 atoms P
formula unit Ca 3 (PO 4 ) 2
c. 4.24 × 1021 formula units Na2HPO4 ×
1 atom P
= 4.24 × 1021 atoms P
formula unit Na 2 HPO 4
Molar mass of CCl2F2 = 12.01 + 2(35.45) + 2(19.00) = 120.91 g/mol
5.56 mg CCl2F2 ×
1g
1 mol
6.022 × 10 23 molecules
×
×
1000 mg 120.91 g
mol
= 2.77 × 1019 molecules CCl2F2
54
CHAPTER 3
5.56 × 10−3 g CCl2F2 ×
STOICHIOMETRY
1 mol CCl 2 F2
2 mol Cl
35.45 g Cl
×
×
120.91 g
1 mol CCl 2 F
mol Cl
= 3.26 × 10−3 g = 3.26 mg Cl
64.
The •2H2O is part of the formula of bauxite (they are called waters of hydration). Combining
elements together, the chemical formula for bauxite would be Al2O5H4.
a. Molar mass = 2(26.98) + 5(16.00) + 4(1.008) = 137.99 g/mol
b. 0.58 mol Al2O3•2H2O ×
26.98 g Al
2 mol Al
×
= 31 g Al
mol Al 2 O 3 • 2H 2 O
mol Al
c. 0.58 mol Al2O3•2H2O ×
6.022 × 10 23 atoms
2 mol Al
×
mol Al
mol Al 2 O 3 • 2H 2 O
= 7.0 × 1023 atoms Al
d. 2.1 × 1024 formula units Al2O3•2H2O ×
1 mol Al2 O 3 • 2H 2 O
6.022 × 10
23
formula units
×
137.99 g
mol
= 480 g Al2O3•2H2O
65.
a. 150.0 g Fe2O3 ×
b. 10.0 mg NO2 ×
1 mol
= 0.9393 mol Fe2O3
159.70 g
1g
1 mol
×
= 2.17 × 10 −4 mol NO2
1000 mg 46.01 g
c. 1.5 × 1016 molecules BF3 ×
66.
a. 20.0 mg C8H10N4O2 ×
1 mol
6.02 × 10 23 molecules
= 2.5 × 10 −8 mol BF3
1g
1 mol
×
= 1.03 × 10 −4 mol C8H10N4O2
1000 mg 194.20 g
b. 2.72 × 1021 molecules C2H5OH ×
1 mol
6.022 × 10 23 molecules
= 4.52 × 10 −3 mol C2H5OH
c. 1.50 g CO2 ×
67.
1 mol
= 3.41 × 10 −2 mol CO2
44.01 g
a. A chemical formula gives atom ratios as well as mole ratios. We will use both ideas to
show how these conversion factors can be used.
Molar mass of C2H5O2N = 2(12.01) + 5(1.008) + 2(16.00) + 14.0l = 75.07 g/mol
CHAPTER 3
STOICHIOMETRY
5.00 g C2H5O2N ×
55
1 mol C 2 H 5O 2 N
6.022 × 10 23 molecules C 2 H 5O 2 N
×
75.07 g C 2 H 5O 2 N
mol C 2 H 5O 2 N
×
1 atom N
= 4.01 × 1022 atoms N
molecule C 2 H 5O 2 N
b. Molar mass of Mg3N2 = 3(24.31) + 2(14.01) = 100.95 g/mol
5.00 g Mg3N2 ×
1 mol Mg 3 N 2
6.022 × 10 23 formula units Mg 3 N 2
×
100.95 g Mg 3 N 2
mol Mg 3 N 2
×
2 atoms N
= 5.97 × 1022 atoms N
mol Mg 3 N 2
c. Molar mass of Ca(NO3)2 = 40.08 + 2(14.01) + 6(16.00) = 164.10 g/mol
5.00 g Ca(NO3)2 ×
1 mol Ca ( NO 3 ) 2
2 mol N
6.022 × 10 23 atoms N
×
×
164.10 g Ca ( NO 3 ) 2
mol Ca ( NO 3 ) 2
mol N
= 3.67 × 1022 atoms N
d. Molar mass of N2O4 = 2(14.01) + 4(16.00) = 92.02 g/mol
5.00 g N2O4 ×
68.
4.24 g C6H6 ×
1 mol N 2 O 4
2 mol N
6.022 × 10 23 atoms N
×
×
92.02 g N 2 O 4
mol N 2 O 4
mol N
= 6.54 × 1022 atoms N
1 mol
= 5.43 × 10 −2 mol C6H6
78.11 g
5.43 × 10 −2 mol C6H6 ×
6.022 × 10 23 molecules
= 3.27 × 1022 molecules C6H6
mol
Each molecule of C6H6 contains 6 atoms C + 6 atoms H = 12 atoms total.
3.27 × 1022 molecules C6H6 ×
12 atoms total
= 3.92 × 1023 atoms total
molecule
0.224 mol H2O ×
18.02 g
= 4.04 g H2O
mol
0.224 mol H2O ×
6.022 × 10 23 molecules
= 1.35 × 1023 molecules H2O
mol
1.35 × 1023 molecules H2O ×
3 atoms total
= 4.05 × 1023 atoms total
molecule
56
CHAPTER 3
2.71 × 1022 molecules CO2 ×
4.50 × 10 −2 mol CO2 ×
1 mol
= 4.50 × 10 −2 mol CO2
6.022 × 10 23 molecules
44.01 g
= 1.98 g CO2
mol
2.71 × 1022 molecules CO2 ×
3.35 × 1022 atoms total ×
3 atoms total
= 8.13 × 1022 atoms total
molecule CO 2
1 molecule
= 5.58 × 1021 molecules CH3OH
6 atoms total
5.58 × 1021 molecules CH3OH ×
9.27 × 10 −3 mol CH3OH ×
69.
32.04 g
= 0.297 g CH3OH
mol
1g
1 mol
×
= 2.839 × 10 −3 mol C6H8O6
1000 mg 176.12 g
2.839 × 10−3 mol ×
6.022 × 10 23 molecules
= 1.710 × 1021 molecules C6H8O6
mol
a. 9(12.01) + 8(1.008) + 4(16.00) = 180.15 g/mol
b. 500. mg ×
1g
1 mol
×
= 2.78 × 10 −3 mol C9H8O4
1000 mg 180.15 g
2.78 × 10 −3 mol ×
71.
1 mol
= 9.27 × 10 −3 mol CH3OH
6.022 × 10 23 molecules
Molar mass of C6H8O6 = 6(12.01) + 8(1.008) + 6(16.00) = 176.12 g/mol
500.0 mg ×
70.
STOICHIOMETRY
a.
2(12.01) + 3(1.008) + 3(35.45) + 2(16.00) = 165.39 g/mol
b. 500.0 g ×
c.
6.022 × 10 23 molecules
= 1.67 × 1021 molecules C9H8O4
mol
1 mol
= 3.023 mol C2H3Cl3O2
165.39 g
2.0 × 10-2 mol ×
165.39 g
= 3.3 g C2H3Cl3O2
mol
d. 5.0 g C2H3Cl3O2 ×
1 mol
6.022 × 10 23 molecules 3 atoms Cl
×
×
molecule
165.39 g
mol
= 5.5 × 1022 atoms of chlorine
CHAPTER 3
72.
STOICHIOMETRY
57
1 mol Cl 1 mol C 2 H 3Cl 3O 2 165.39 g C 2 H 3Cl 3O 2
×
×
= 1.6 g chloral hydrate
35.45 g
3 mol Cl
mol C 2 H 3Cl 3O 2
e.
1.0 g Cl ×
f.
500 molecules ×
1 mol
6.022 × 10
23
molecules
×
165.39 g
= 1.373 × 10−19 g C2H3Cl3O2
mol
As we shall see in later chapters, the formula written as (CH3)2N2O tries to tell us something
about how the atoms are attached to each other. For our purposes in this problem, we can
write the formula as C2H6N2O.
a. 2(12.01) + 6(1.008) + 2(14.01) + 1(16.00) = 74.09 g/mol
b.
250 mg ×
1g
1 mol
×
= 3.4 × 10−3 mol
1000 mg 74.09 g
d. 1.0 mol C2H6N2O ×
c.
0.050 mol ×
74.09 g
= 3.7 g
mol
6.022 × 10 23 molecules C 2 H 6 N 2 O
6 atoms of H
×
mol C 2 H 6 N 2 O
molecule C 2 H 6 N 2 O
= 3.6 × 1024 atoms of hydrogen
e. 1.0 × 106 molecules ×
f.
1 mol
6.022 × 10
23
molecules
1 mol
1 molecule ×
6.022 × 10
23
molecules
×
×
74.09 g
= 1.2 × 10−16 g
mol
74.09 g
= 1.230 × 10−22 g C2H6N2O
mol
Percent Composition
73.
a. C3H4O2: Molar mass = 3(12.01) + 4(1.008) + 2(16.00) = 36.03 + 4.032 + 32.00
= 72.06 g/mol
36.03 g C
Mass % C =
× 100 = 50.00% C
72.06 g compound
Mass % H =
4.032 g H
× 100 = 5.595% H
72.06 g compound
Mass % O = 100.00 − (50.00 + 5.595) = 44.41% O or:
%O=
32.00 g
× 100 = 44.41% O
72.06 g
b. C4H6O2: Molar mass = 4(12.01) + 6(1.008) + 2(16.00) = 48.04 + 6.048 + 32.00
= 86.09 g/mol
Mass % C =
48.04 g
6.048 g
× 100 = 55.80% C; mass % H =
× 100 = 7.025% H
86.09 g
86.09 g
Mass % O = 100.00 − (55.80 + 7.025) = 37.18% O
58
CHAPTER 3
STOICHIOMETRY
c. C3H3N: Molar mass = 3(12.01) + 3(1.008) + 1(14.01) = 36.03 + 3.024 + 14.01
= 53.06 g/mol
74.
Mass % C =
36.03 g
3.024 g
× 100 = 67.90% C; mass % H =
× 100 = 5.699% H
53.06 g
53.06 g
Mass % N =
14.01 g
× 100 = 26.40% N or % N = 100.00 − (67.90 + 5.699)
53.06 g
= 26.40% N
In 1 mole of YBa2Cu3O7, there are 1 mole of Y, 2 moles of Ba, 3 moles of Cu, and 7 moles
of O.
 137.3 g Ba 
 88.91 g Y 

 + 2 mol Ba 
Molar mass = 1 mol Y 
 mol Ba 
 mol Y 
 63.55 g Cu 
 + 7 mol O
+ 3 mol Cu 
 mol Cu 
 16.00 g O 


 mol O 
Molar mass = 88.91 + 274.6 + 190.65 + 112.00 = 666.2 g/mol
Mass % Y =
Mass % Cu =
75.
88.91 g
274.6 g
× 100 = 13.35% Y; mass % Ba =
× 100 = 41.22% Ba
666.2 g
666.2 g
112.0 g
190.65 g
× 100 = 28.62% Cu; mass % O =
× 100 = 16.81% O
666.2 g
666.2 g
14.01 g N
× 100 = 46.68% N
30.01 g NO
NO: Mass % N =
NO2: Mass % N =
14.01 g N
× 100 = 30.45% N
46.01 g NO 2
N2O: Mass % N =
2(14.01) g N
× 100 = 63.65% N
44.02 g N 2 O
From the calculated mass percents, only NO is 46.7% N by mass, so NO could be this
species. Any other compound having NO as an empirical formula could also be the
compound.
76.
a.
C8H10N4O2: Molar mass = 8(12.01) + 10(1.008) + 4(14.0l) + 2(16.00) = 194.20 g/mol
Mass % C =
96.08 g
8(12.01) g C
× 100 =
× 100 = 49.47% C
194.20 g
194.20 g C8 H10 N 4 O 2
b. C12 H22O11: Molar mass = 12(12.01) + 22(1.008) + 11(16.00) = 342.30 g/mol
Mass % C =
12(12.01) g C
× 100 = 42.10% C
342.30 g C12 H 22 O11
CHAPTER 3
c.
STOICHIOMETRY
59
C2H5OH: Molar mass = 2(12.01) + 6(1.008) + 1(16.00) = 46.07 g/mol
Mass % C =
2(12.01) g C
× 100 = 52.14% C
46.07 g C 2 H 5 OH
The order from lowest to highest mass percentage of carbon is:
sucrose (C12H22O11) < caffeine (C8H10N4O2) < ethanol (C2H5OH)
77.
There are 0.390 g Cu for every 100.000 g of fungal laccase. Assuming 100.00 g fungal
laccase:
1 mol Cu
1 mol fungal laccase
×
Mol fungal laccase = 0.390 g Cu ×
= 1.53 × 10 −3 mol
63.55 g Cu
4 mol Cu
x g fungal laccase
100.000 g
=
, x = molar mass = 6.54 × 104 g/mol
−
3
mol fungal laccase 1.53 × 10 mol
78.
There are 0.347 g Fe for every 100.000 g hemoglobin (Hb). Assuming 100.000 g
hemoglobin:
Mol Hb = 0.347 g Fe ×
1 mol Fe
1 mol Hb
×
= 1.55 × 10−3 mol Hb
55.85 g Fe 4 mol Fe
x g Hb
100.000 g Hb
=
, x = molar mass = 6.45 × 104 g/mol
−3
mol Hb
1.55 × 10 mol Hb
Empirical and Molecular Formulas
79.
 1.008 g H 
 12.01 g C 

 + 2 mol H 
a. Molar mass of CH2O = 1 mol C 
 mol H 
 mol C 
 16.00 g O 
 = 30.03 g/mol
+ 1 mol O 
 mol O 
%C=
12.01 g C
2.016 g H
× 100 = 39.99% C; % H =
× 100 = 6.713% H
30.03 g CH 2 O
30.03 g CH 2 O
%O=
16.00 g O
× 100 = 53.28% O or % O = 100.00 − (39.99 + 6.713) = 53.30%
30.03 g CH 2 O
b. Molar mass of C6H12O6 = 6(12.01) + 12(1.008) + 6(16.00) = 180.16 g/mol
%C=
76.06 g C
× 100 = 40.00%;
180.16 g C 6 H 12 O 6
% O = 100.00 − (40.00 + 6.714) = 53.29%
%H=
12.(1.008) g
× 100 = 6.714%
180.16 g
60
CHAPTER 3
STOICHIOMETRY
c. Molar mass of HC2H3O2 = 2(12.01) + 4(1.008) + 2(16.00) = 60.05 g/mol
%C=
24.02 g
× 100 = 40.00%;
60.05 g
%H=
4.032 g
× 100 = 6.714%
60.05 g
% O = 100.00 − (40.00 + 6.714) = 53.29%
80.
All three compounds have the same empirical formula, CH2O, and different molecular
formulas. The composition of all three in mass percent is also the same (within rounding
differences). Therefore, elemental analysis will give us only the empirical formula.
81.
a. The molecular formula is N2O4. The smallest whole number ratio of the atoms (the
empirical formula) is NO2.
b. Molecular formula: C3H6; empirical formula: CH2
c. Molecular formula: P4O10; empirical formula: P2O5
d. Molecular formula: C6H12O6; empirical formula: CH2O
82.
a. SNH: Empirical formula mass = 32.07 + 14.01 + 1.008 = 47.09 g/mol
188.35 g
= 4.000; so the molecular formula is (SNH)4 or S4N4H4.
47.09 g
b. NPCl2: Empirical formula mass = 14.01 + 30.97 + 2(35.45) = 115.88 g/mol
347.64 g
= 3.0000; molecular formula is (NPCl2)3 or N3P3Cl6.
115.88 g
c. CoC4O4: 58.93 + 4(12.01) + 4(16.00) = 170.97 g/mol
341.94 g
= 2.0000; molecular formula: Co2C8O8
170.97 g
d. SN: 32.07 + 14.01 = 46.08 g/mol;
83.
184.32 g
= 4.000; molecular formula: S4N4
46.08 g
Out of 100.00 g of compound, there are:
48.64 g C ×
1 mol C
1 mol H
= 4.050 mol C; 8.16 g H ×
= 8.10 mol H
12.01 g C
1.008 g H
% O = 100.00 – 48.64 – 8.16 = 43.20%; 43.20 g O ×
Dividing each mole value by the smallest number:
1 mol O
= 2.700 mol O
16.00 g O
CHAPTER 3
STOICHIOMETRY
61
4.050
2.700
8.10
= 1.500;
= 3.00;
= 1.000
2.700
2.700
2.700
Because a whole number ratio is required, the C : H : O ratio is 1.5 : 3 : 1 or 3 : 6 : 2. So the
empirical formula is C3H6O2.
84.
Assuming 100.00 g of nylon-6:
63.68 g C ×
9.80 g H ×
1 mol C
1 mol N
= 5.302 mol C; 12.38 g N ×
= 0.8837 mol N
12.01 g C
14.01 g N
1 mol H
1 mol O
= 9.72 mol H; 14.14 g O ×
= 0.8838 mol O
1.008 g H
16.00 g O
Dividing each mole value by the smallest number:
0.8838
9.72
5.302
= 6.000;
= 11.0;
= 1.000
0.8837
0.8837
0.8837
The empirical formula for nylon-6 is C6H11NO
85.
Compound I: Mass O = 0.6498 g HgxOy − 0.6018 g Hg = 0.0480 g O
0.6018 g Hg ×
0.0480 g O ×
1 mol Hg
= 3.000 × 10−3 mol Hg
200.6 g Hg
1 mol O
= 3.00 × 10−3 mol O
16.00 g O
The mole ratio between Hg and O is 1 : 1, so the empirical formula of compound I is HgO.
Compound II: Mass Hg = 0.4172 g HgxOy − 0.016 g O = 0.401 g Hg
0.401 g Hg ×
1 mol Hg
1 mol O
= 2.00 × 10−3 mol Hg; 0.016 g O ×
= 1.0 × 10−3 mol O
16.00 g O
200.6 g Hg
The mole ratio between Hg and O is 2 : 1, so the empirical formula is Hg2O.
86.
1.121 g N ×
1 mol N
14.01 g N
= 8.001 × 10 −2 mol N; 0.161 g H ×
0.480 g C ×
1 mol C
12.01 g C
= 4.00 × 10 −2 mol C; 0.640 g O ×
Dividing all mole values by the smallest number:
1 mol H
1.008 g H
1 mol O
16.00 g O
= 1.60 × 10 −1 mol H
= 4.00 × 10 −2 mol O
62
CHAPTER 3
8.001 × 10 −2
4.00 × 10 − 2
1.60 × 10 −1
= 2.00;
4.00 × 10 − 2
= 4.00;
4.00 × 10 −2
4.00 × 10 − 2
STOICHIOMETRY
= 1.00
The empirical formula is N2H4CO.
87.
Out of 100.0 g, there are:
69.6 g S ×
1 mol S
32.07 g S
= 2.17 mol S; 30.4 g N ×
1 mol N
14.01 g N
= 2.17 mol N
The empirical formula is SN because the mole values are in a 1 : 1 mole ratio.
The empirical formula mass of SN is ~ 46 g/mol. Because 184/46 = 4.0, the molecular
formula is S4N4.
88.
Assuming 100.0 g of compound:
26.7 g P ×
1 mol P
30.97 g P
61.2 g Cl ×
= 0.862 mol P; 12.1 g N ×
1 mol Cl
35.45 g Cl
1 mol N
14.01 g N
= 0.864 mol N
= 1.73 mol Cl
1.73
= 2.01; the empirical formula is PNCl2.
0.862
The empirical formula mass is ≈ 31.0 + 14.0 + 2(35.5) = 116 g/mol.
Molar mass
Empirical formula mass
89.
=
580
= 5.0; the molecular formula is (PNCl2)5 = P5N5Cl10.
116
Assuming 100.00 g of compound:
47.08 g C ×
46.33 g Cl ×
1 mol C
1 mol H
= 3.920 mol C; 6.59 g H ×
= 6.54 mol H
1.008 g H
12.01 g C
1 mol Cl
= 1.307 mol Cl
35.45 g Cl
Dividing all mole values by 1.307 gives:
1.307
3.920
6.54
= 2.999;
= 5.00;
= 1.000
1.307
1.307
1.307
The empirical formula is C3H5Cl.
The empirical formula mass is 3(12.01) + 5(1.008) + 1(35.45) = 76.52 g/mol.
CHAPTER 3
STOICHIOMETRY
63
Molar mass
153
=
= 2.00 ; the molecular formula is (C3H5Cl)2 = C6H10Cl2.
Empirical formula mass
76.52
90.
Assuming 100.00 g of compound (mass oxygen = 100.00 g − 41.39 g C − 3.47 g H
= 55.14 g O):
41.39 g C ×
1 mol C
12.01 g C
= 3.446 mol C; 3.47 g H ×
55.14 g O ×
1 mol O
16.00 g O
= 3.446 mol O
1 mol H
1.008 g H
= 3.44 mol H
All are the same mole values, so the empirical formula is CHO. The empirical formula mass
is 12.01 + 1.008 + 16.00 = 29.02 g/mol.
Molar mass =
15.0 g
0.129 mol
= 116 g/mol
Molar mass
116
=
= 4.00; molecular formula = (CHO)4 = C4H4O4
Empirical mass
29.02
91.
When combustion data are given, it is assumed that all the carbon in the compound ends up
as carbon in CO2 and all the hydrogen in the compound ends up as hydrogen in H2O. In the
sample of fructose combusted, the masses of C and H are:
mass C = 2.20 g CO2 ×
1 mol CO 2
44.01 g CO 2
mass H = 0.900 g H2O ×
×
1 mol C
12.01 g C
×
= 0.600 g C
mol CO 2
mol C
1 mol H 2 O
2 mol H
1.008 g H
×
×
= 0.101 g H
mol H 2 O
mol H
18.02 g H 2 O
Mass O = 1.50 g fructose − 0.600 g C − 0.101 g H = 0.799 g O
So, in 1.50 g of the fructose, we have:
0.600 g C ×
1 mol H
1 mol C
= 0.0500 mol C; 0.101 g H ×
= 0.100 mol H
1.008 g H
12.01 g C
0.799 g O ×
1 mol O
= 0.0499 mol O
16.00 g O
Dividing by the smallest number:
92.
0.100
= 2.00; the empirical formula is CH2O.
0.0499
This compound contains nitrogen, and one way to determine the amount of nitrogen in the
compound is to calculate composition by mass percent. We assume that all the carbon in
33.5 mg CO2 came from the 35.0 mg of compound and all the hydrogen in 41.1 mg H2O
came from the 35.0 mg of compound.
64
CHAPTER 3
3.35 × 10 −2 g CO2 ×
Mass % C =
×
1 mol C
12.01 g C
×
= 9.14 × 10 −3 g C
mol CO 2
mol C
9.14 × 10 −3 g C
3.50 × 10 − 2 g compound
4.11 × 10 −2 g H2O ×
Mass % H =
1 mol CO 2
44.01 g CO 2
STOICHIOMETRY
× 100 = 26.1% C
1 mol H 2 O
2 mol H
1.008 g H
×
×
= 4.60 × 10 −3 g H
18.02 g H 2 O
mol H 2 O
mol H
4.60 × 10 −3 g H
3.50 × 10 − 2 g compound
× 100 = 13.1% H
The mass percent of nitrogen is obtained by difference:
Mass % N = 100.0 − (26.1 + 13.1) = 60.8% N
Now perform the empirical formula determination by first assuming 100.0 g of compound.
Out of 100.0 g of compound, there are:
26.1 g C ×
1 mol C
1 mol H
= 2.17 mol C; 13.1 g H ×
= 13.0 mol H
1.008 g H
12.01 g C
60.8 g N ×
1 mol N
= 4.34 mol N
14.01 g N
Dividing all mole values by 2.17 gives:
2.17
13.0
4.34
= 1.00;
= 5.99;
= 2.00
2.17
2.17
2.17
The empirical formula is CH6N2.
93.
The combustion data allow determination of the amount of hydrogen in cumene. One way to
determine the amount of carbon in cumene is to determine the mass percent of hydrogen in
the compound from the data in the problem; then determine the mass percent of carbon by
difference (100.0 − mass % H = mass % C).
42.8 mg H2O ×
Mass % H =
1g
1000 mg
×
2.016 g H
1000 mg
×
= 4.79 mg H
18.02 g H 2 O
g
4.79 mg H
× 100 = 10.1% H; mass % C = 100.0 − 10.1 = 89.9% C
47.6 mg cumene
Now solve the empirical formula problem. Out of 100.0 g cumene, we have:
89.9 g C ×
1 mol C
1 mol H
= 7.49 mol C; 10.1 g H ×
= 10.0 mol H
12.01 g C
1.008 g H
CHAPTER 3
STOICHIOMETRY
65
4
10.0
= 1.34 ≈ ; the mole H to mole C ratio is 4 : 3. The empirical formula is C3H4.
7.49
3
Empirical formula mass ≈ 3(12) + 4(1) = 40 g/mol.
The molecular formula must be (C3H4)3 or C9H12 because the molar mass of this formula will
be between 115 and 125 g/mol (molar mass ≈ 3 × 40 g/mol = 120 g/mol).
94.
There are several ways to do this problem. We will determine composition by mass percent:
16.01 mg CO2 ×
%C=
×
12.01 g C
1000 mg
×
= 4.369 mg C
44.01 g CO 2
g
4.369 mg C
× 100 = 40.91% C
10.68 mg compound
4.37 mg H2O ×
%H=
1g
1000 mg
1g
1000 mg
×
1000 mg
2.016 g H
×
= 0.489 mg H
18.02 g H 2 O
g
0.489 mg
× 100 = 4.58% H; % O = 100.00 − (40.91 + 4.58) = 54.51% O
10.68 mg
So, in 100.00 g of the compound, we have:
40.91 g C ×
1 mol H
1 mol C
= 3.406 mol C; 4.58 g H ×
= 4.54 mol H
1.008 g H
12.01 g C
54.51 g O ×
1 mol O
= 3.407 mol O
16.00 g O
Dividing by the smallest number:
4.54
4
= 1.33 ≈ ; the empirical formula is C3H4O3.
3.406
3
The empirical formula mass of C3H4O3 is ≈ 3(12) + 4(1) + 3(16) = 88 g/mol.
Because
176.1
= 2.0, the molecular formula is C6H8O6.
88
Balancing Chemical Equations
95.
When balancing reactions, start with elements that appear in only one of the reactants and one
of the products, and then go on to balance the remaining elements.
a. C6H12O6(s) + O2(g) → CO2(g) + H2O(g)
Balance C atoms: C6H12O6 + O2 → 6 CO2 + H2O
Balance H atoms: C6H12O6 + O2 → 6 CO2 + 6 H2O
Lastly, balance O atoms: C6H12O6(s) + 6 O2(g) → 6 CO2(g) + 6 H2O(g)
66
CHAPTER 3
STOICHIOMETRY
b. Fe2S3(s) + HCl(g) → FeCl3(s) + H2S(g)
Balance Fe atoms: Fe2S3 + HCl → 2 FeCl3 + H2S
Balance S atoms: Fe2S 3 + HCl → 2 FeCl3 + 3 H2S
There are 6 H and 6 Cl on right, so balance with 6 HCl on left:
Fe2S3(s) + 6 HCl(g) → 2 FeCl3(s) + 3 H2S(g).
c. CS2(l) + NH3(g) → H2S(g) + NH4SCN(s)
C and S balanced; balance N:
CS2 + 2 NH3 → H2S + NH4SCN
H is also balanced. CS2(l) + 2 NH3(g) → H2S(g) + NH4SCN(s)
96.
An important part to this problem is writing out correct formulas. If the formulas are
incorrect, then the balanced reaction is incorrect.
a. C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(g)
b. 3 Pb(NO3)2(aq) + 2 Na3PO4(aq) → Pb3(PO4)2(s) + 6 NaNO3(aq)
c. Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g)
d. Sr(OH)2(aq) + 2 HBr(aq) → 2H2O(l) + SrBr2(aq)
MnO2
catalyst
97.
2 H2O2(aq)
98.
Fe3O4(s) + 4 H2(g) → 3 Fe(s) + 4 H2O(g)
2 H2O(l) + O2(g)
Fe3O4(s) + 4 CO(g) → 3 Fe(s) + 4 CO2(g)
99.
a. 3 Ca(OH)2(aq) + 2 H3PO4(aq) → 6 H2O(l) + Ca3(PO4)2(s)
b. Al(OH)3(s) + 3 HCl(aq) → AlCl3(aq) + 3 H2O(l)
c. 2 AgNO3(aq) + H2SO4(aq) → Ag2SO4(s) + 2 HNO3(aq)
100.
a. 2 KO2(s) + 2 H2O(l) → 2 KOH(aq) + O2(g) + H2O2(aq) or
4 KO2(s) + 6 H2O(l) → 4 KOH(aq) + O2(g) + 4 H2O2(aq)
b. Fe2O3(s) + 6 HNO3(aq) → 2 Fe(NO3)3(aq) + 3 H2O(l)
c. 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g)
CHAPTER 3
STOICHIOMETRY
67
d. PCl5(l) + 4 H2O(l) → H3PO4(aq) + 5 HCl(g)
e. 2 CaO(s) + 5 C(s) → 2 CaC2(s) + CO2(g)
f.
2 MoS2(s) + 7 O2(g) → 2 MoO3(s) + 4 SO2(g)
g. FeCO3(s) + H2CO3(aq) → Fe(HCO3)2(aq)
101.
a. The formulas of the reactants and products are C6H6(l) + O2(g) → CO2(g) + H2O(g). To
balance this combustion reaction, notice that all of the carbon in C6H6 has to end up as
carbon in CO2 and all of the hydrogen in C6H6 has to end up as hydrogen in H2O. To
balance C and H, we need 6 CO2 molecules and 3 H2O molecules for every 1 molecule of
C6H6. We do oxygen last. Because we have 15 oxygen atoms in 6 CO2 molecules and 3
H2O molecules, we need 15/2 O2 molecules in order to have 15 oxygen atoms on the
reactant side.
15
C6H6(l) +
2
O2(g) → 6 CO2(g) + 3 H2O(g); multiply by two to give whole numbers.
2 C6H6(l) + 15 O2(g) → 12 CO2(g) + 6 H2O(g)
b. The formulas of the reactants and products are C4H10(g) + O2(g) → CO2(g) + H2O(g).
C4H10(g) +
13
2
O2(g) → 4 CO2(g) + 5 H2O(g); multiply by two to give whole numbers.
2 C4H10(g) + 13 O2(g) → 8 CO2(g) + 10 H2O(g)
c. C12H22O11(s) + 12 O2(g) → 12 CO2(g) + 11 H2O(g)
d. 2 Fe(s) +
3
2
e. 2 FeO(s) +
O2(g) → Fe2O3(s); for whole numbers: 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s)
1
2
O2(g) → Fe2O3(s); for whole numbers, multiply by two.
4 FeO(s) + O2(g) → 2 Fe2O3(s)
102.
a. 16 Cr(s) + 3 S8(s) → 8 Cr2S3(s)
b. 2 NaHCO3(s) → Na2CO3(s) + CO2(g) + H2O(g)
c. 2 KClO3(s) → 2 KCl(s) + 3 O2(g)
d. 2 Eu(s) + 6 HF(g) → 2 EuF3(s) + 3 H2(g)
103.
a. SiO2(s) + C(s) → Si(s) + CO(g); Si is balanced.
Balance oxygen atoms: SiO2 + C → Si + 2 CO
Balance carbon atoms: SiO2(s) + 2 C(s) → Si(s) + 2 CO(g)
68
CHAPTER 3
STOICHIOMETRY
b. SiCl4(l) + Mg(s) → Si(s) + MgCl2(s); Si is balanced.
Balance Cl atoms: SiCl4 + Mg → Si + 2 MgCl2
Balance Mg atoms: SiCl4(l) + 2 Mg(s) → Si(s) + 2 MgCl2(s)
c. Na2SiF6(s) + Na(s) → Si(s) + NaF(s); Si is balanced.
Balance F atoms:
Na2SiF6 + Na → Si + 6 NaF
Balance Na atoms: Na2SiF6(s) + 4 Na(s) → Si(s) + 6 NaF(s)
104.
CaSiO3(s) + 6 HF(aq) → CaF2(aq) + SiF4(g) + 3 H2O(l)
Reaction Stoichiometry
105.
The stepwise method to solve stoichiometry problems is outlined in the text. Instead of
calculating intermediate answers for each step, we will combine conversion factors into one
calculation. This practice reduces round-off error and saves time.
Fe2O3(s) + 2 Al(s) → 2 Fe(l) + Al2O3(s)
15.0 g Fe ×
106.
1 mol Fe
2 mol Al 26.98 g Al
×
= 0.269 mol Fe; 0.269 mol Fe ×
= 7.26 g Al
55.85 g Fe
2 mol Fe
mol Al
0.269 mol Fe ×
1 mol Fe 2 O 3 159.70 g Fe 2 O 3
×
= 21.5 g Fe2O3
2 mol Fe
mol Fe 2 O 3
0.269 mol Fe ×
1 mol Al2 O 3 101.96 g Al2 O 3
×
= 13.7 g Al2O3
2 mol Fe
mol Al2 O 3
10 KClO3(s) + 3 P4(s) → 3 P4O10(s) + 10 KCl(s)
52.9 g KClO3 ×
107.
1.000 kg Al ×
1 mol KClO3
3 mol P4 O10
283.88 g P4 O10
×
×
= 36.8 g P4O10
122.55 g KClO3 10 mol KClO3
mol P4 O10
3 mol NH 4 ClO 4
117.49 g NH 4 ClO 4
1000 g Al
1 mol Al
×
×
×
kg Al
26.98 g Al
3 mol Al
mol NH 4 ClO 4
= 4355 g = 4.355 kg NH4ClO4
108.
a. Ba(OH)2•8H2O(s) + 2 NH4SCN(s) → Ba(SCN)2(s) + 10 H2O(l) + 2 NH3(g)
b. 6.5 g Ba(OH)2•8H2O ×
1 mol Ba(OH) 2 • 8H 2 O
= 0.0206 mol = 0.021 mol
315.4 g
CHAPTER 3
STOICHIOMETRY
0.021 mol Ba(OH)2•8H2O ×
69
2 mol NH 4 SCN
76.13 g NH 4 SCN
×
1 mol Ba(OH) 2 • 8H 2 O
mol NH 4 SCN
= 3.2 g NH4SCN
109.
a. 1.0 × 102 mg NaHCO3 ×
1 mol NaHCO3
1 mol C 6 H 8O 7
1g
×
×
1000 mg 84.01 g NaHCO3 3 mol NaHCO3
×
b. 0.10 g NaHCO3 ×
192.12 g C 6 H 8O 7
= 0.076 g or 76 mg C6H8O7
mol C 6 H 8O 7
1 mol NaHCO3
3 mol CO 2
44.01 g CO 2
×
×
84.01 g NaHCO3 3 mol NaHCO3
mol CO 2
= 0.052 g or 52 mg CO2
110.
a.
1.00 × 102 g C7H6O3 ×
1 mol C 7 H 6 O 3
1 mol C 4 H 6 O 3
102.09 g C 4 H 6 O 3
×
×
138.12 g C 7 H 6 O 3
1 mol C 7 H 6 O 3
1 mol C 4 H 6 O 3
= 73.9 g C4H6O3
b.
1.00 × 102 g C7H6O3 ×
1 mol C 7 H 6 O 3
1 mol C9 H 8O 4
180.15 g C9 H 8O 4
×
×
138.12 g C 7 H 6 O 3
1 mol C 7 H 6 O 3
mol C9 H 8O 4
= 1.30 × 102 g aspirin
111.
1.0 × 104 kg waste ×
1 mol C5 H 7 O 2 N
3.0 kg NH 4 +
1 mol NH 4 +
1000 g
×
×
×
+
100 kg waste
kg
18.04 g NH 4
55 mol NH 4 +
×
113.12 g C 5 H 7 O 2 N
= 3.4 × 104 g tissue if all NH4+ converted
mol C 5 H 7 O 2 N
Because only 95% of the NH4+ ions react:
mass of tissue = (0.95)(3.4 × 104 g) = 3.2 × 104 g or 32 kg bacterial tissue
112.
1.0 × 103 g phosphorite ×
75 g Ca 3 (PO 4 ) 2
1 mol Ca 3 (PO 4 ) 2
×
100 g phosphorite
310.18 g Ca 3 (PO 4 ) 2
×
113.
1.0 ton CuO ×
1 mol P4
2 mol Ca 3 (PO 4 ) 2
×
123.88 g P4
= 150 g P4
mol P4
907 kg 1000 g
1 mol CuO
1 mol C
12.01 g C 100. g coke
×
×
×
×
×
ton
kg
79.55 g CuO 2 mol CuO
mol C
95 g C
= 7.2 × 104 g or 72 kg coke
114.
2 LiOH(s) + CO2(g) → Li2CO3(aq) + H2O(l)
The total volume of air exhaled each minute for the 7 astronauts is 7 × 20. = 140 L/min.
70
CHAPTER 3
25,000 g LiOH ×
STOICHIOMETRY
1 mol CO 2
44.01 g CO 2
1 mol LiOH
100 g air
×
×
×
23.95 g LiOH
2 mol LiOH
mol CO 2
4.0 g CO 2
×
1 mL air
1L
1 min
1h
×
×
×
= 68 h = 2.8 days
0.0010 g air
1000 mL
140 L air
60 min
Limiting Reactants and Percent Yield
115.
The product formed in the reaction is NO2; the other species present in the product representtation is excess O2. Therefore, NO is the limiting reactant. In the pictures, 6 NO molecules
react with 3 O2 molecules to form 6 NO2 molecules.
6 NO(g) + 3 O2(g) → 6 NO2(g)
For smallest whole numbers, the balanced reaction is:
2 NO(g) + O2(g) → 2 NO2(g)
116.
In the following table we have listed three rows of information. The “Initial” row is the
number of molecules present initially, the “Change” row is the number of molecules that
react to reach completion, and the “Final” row is the number of molecules present at
completion. To determine the limiting reactant, let’s calculate how much of one reactant is
necessary to react with the other.
10 molecules O2 ×
4 molecules NH 3
= 8 molecules NH3 to react with all of the O2
5 molecules O 2
Because we have 10 molecules of NH3 and only 8 molecules of NH3 are necessary to react
with all of the O2, O2 is limiting. Now use the 10 molecules of O2 and the molecule
relationships given in the balanced equation to determine the number of molecules of each
product formed, then complete the table.
4 NH3(g)
Initial
Change
Final
+
10 molecules
−8 molecules
2 molecules
5 O2(g)
10 molecules
−10 molecules
0
→
4 NO(g)
0
+8 molecules
8 molecules
+
6 H2O(g)
0
+12 molecules
12 molecules
The total number of molecules present after completion = 2 molecules NH3 + 0 molecules O2
+ 8 molecules NO + 12 molecules H2O = 22 molecules.
117.
a. The strategy we will generally use to solve limiting reactant problems is to assume each
reactant is limiting, and then calculate the quantity of product each reactant could
produce if it were limiting. The reactant that produces the smallest quantity of product
is the limiting reactant (runs out first) and therefore determines the mass of product that
can be produced.
Assuming N2 is limiting:
1.00 × 103 g N2 ×
2 mol NH 3
17.03 g NH 3
1 mol N 2
×
×
= 1.22 × 103 g NH3
28.02 g N 2
mol N 2
mol NH 3
CHAPTER 3
STOICHIOMETRY
71
Assuming H2 is limiting:
2 mol NH 3
17.03 g NH 3
1 mol H 2
×
×
= 2.82 × 103 g NH3
2.016 g H 2
3 mol H 2
mol NH 3
5.00 × 102 g H2 ×
Because N2 produces the smaller mass of product (1220 g vs. 2820 g NH3), N2 is limiting
and 1220 g NH3 can be produced. As soon as 1220 g of NH3 is produced, all of the N2 has
run out. Even though we have enough H2 to produce more product, there is no more N2
present as soon as 1220 g of NH3 have been produced.
b. 1.00 × 103 g N2 ×
1 mol N 2
3 mol H 2
2.016 g H 2
×
×
= 216 g H2 reacted
28.02 g N 2
mol N 2
mol H 2
Excess H2 = 500. g H2 initially – 216 g H2 reacted = 284 g H2 in excess (unreacted)
118.
Ca3(PO4)2 + 3 H2SO4 → 3 CaSO4 + 2 H3PO4
Assuming Ca3(PO4)2 is limiting:
1.0 × 103 g Ca3(PO4)2 ×
1 mol Ca 3 (PO 4 ) 2
136.15 g CaSO 4
3 mol CaSO 4
×
×
mol CaSO 4
mol Ca 3 (PO 4 ) 2
310.18 g Ca 3 (PO 4 ) 2
= 1300 g CaSO4
Assuming concentrated H2SO4 reagent is limiting:
1.0 × 103 g conc. H2SO4 ×
98 g H 2SO 4
1 mol H 2SO 4
×
100 g conc. H 2SO 4 98.09 g H 2SO 4
×
3 mol CaSO 4 136.15 g CaSO 4
×
= 1400 g CaSO4
3 mol H 2SO 4
mol CaSO 4
Because Ca3(PO4)2 produces the smaller quantity of product, Ca3(PO4)2 is limiting and
1300 g CaSO4 can be produced.
1.0 × 103 g Ca3(PO4)2 ×
2 mol H 3 PO 4
97.99 g H 3 PO 4
1 mol Ca 3 (PO 4 ) 2
×
×
mol Ca 3 (PO 4 ) 2
mol H 3 PO 4
310.18 g Ca 3 (PO 4 ) 2
= 630 g H3PO4 produced
119.
Assuming BaO2 is limiting:
1.50 g BaO2 ×
1 mol BaO 2
1 mol H 2 O 2 34.02 g H 2 O 2
×
×
= 0.301 g H2O2
169.3 g BaO 2
mol BaO 2
mol H 2 O 2
Assuming HCl is limiting:
25.0 mL ×
0.0272 g HCl 1 mol HCl 1 mol H 2 O 2 34.02 g H 2 O 2
×
×
×
= 0.317 g H2O2
mL
36.46 g HCl 2 mol HCl
mol H 2 O 2
72
CHAPTER 3
STOICHIOMETRY
BaO2 produces the smaller amount of H2O2, so it is limiting and a mass of 0.301 g of H2O2
can be produced.
Initial mol HCl present: 25.0 mL ×
0.0272 g HCl
1 mol HCl
×
= 1.87 × 10−2 mol HCl
mL
36.46 g HCl
The amount of HCl reacted:
1.50 g BaO2 ×
1 mol BaO 2
2 mol HCl
×
= 1.77 × 10 −2 mol HCl
169.3 g BaO 2
mol BaO 2
Excess mol HCl = 1.87 × 10 −2 mol − 1.77 × 10 −2 mol = 1.0 × 10 −3 mol HCl
Mass of excess HCl = 1.0 × 10 −3 mol HCl ×
120.
36.46 g HCl
= 3.6 × 10 −2 g HCl unreacted
mol HCl
Assuming Ag2O is limiting:
25.0 g Ag2O ×
2 mol AgC10 H 9 N 4SO 2 357.18 g AgC10 H 9 N 4SO 2
1 mol Ag 2 O
×
×
mol Ag 2 O
mol AgC10 H 9 N 4SO 2
231.8 g Ag 2 O
= 77.0 g AgC10H9N4SO2
Assuming C10H10N4SO2 is limiting:
50.0 g C10H10N4SO2 ×
1 mol C10 H10 N 4SO 2
2 mol AgC10 H 9 N 4SO 2
×
250.29 g C10 H10 N 4SO 2
2 mol C10 H10 N 4SO 2
×
357.18 g AgC10 H 9 N 4SO 2
= 71.4 g AgC10H9N4SO2
mol AgC10 H 9 N 4SO 2
Because C10H10N4SO2 produces the smaller amount of product, it is limiting and 71.4 g of
silver sulfadiazine can be produced.
121.
To solve limiting-reagent problems, we will generally assume each reactant is limiting and
then calculate how much product could be produced from each reactant. The reactant that
produces the smallest amount of product will run out first and is the limiting reagent.
5.00 × 106 g NH3 ×
5.00 × 106 g O2 ×
5.00 × 106 g CH4 ×
1 mol NH 3
2 mol HCN
×
= 2.94 × 105 mol HCN
17.03 g NH 3
2 mol NH 3
1 mol O 2
2 mol HCN
×
= 1.04 × 105 mol HCN
32.00 g O 2
3 mol O 2
1 mol CH 4
2 mol HCN
×
= 3.12 × 105 mol HCN
16.04 g CH 4
2 mol CH 4
CHAPTER 3
STOICHIOMETRY
73
O2 is limiting because it produces the smallest amount of HCN. Although more product could
be produced from NH3 and CH4, only enough O2 is present to produce 1.04 × 105 mol HCN.
The mass of HCN produced is:
1.04 × 105 mol HCN ×
5.00 × 106 g O2 ×
122.
27.03 g HCN
= 2.81 × 106 g HCN
mol HCN
1 mol O 2
6 mol H 2 O 18.02 g H 2 O
×
×
= 5.63 × 106 g H2O
32.00 g O 2
3 mol O 2
1 mol H 2 O
If C3H6 is limiting:
15.0 g C3H6 ×
1 mol C 3 H 6
2 mol C 3 H 3 N
53.06 g C 3 H 3 N
×
×
= 18.9 g C3H3N
42.08 g C 3 H 6
2 mol C 3 H 6
mol C 3 H 3 N
If NH3 is limiting:
5.00 g NH3 ×
1 mol NH 3
2 mol C 3 H 3 N
53.06 g C 3 H 3 N
×
×
= 15.6 g C3H3N
17.03 g NH 3
2 mol NH 3
mol C 3 H 3 N
If O2 is limiting:
10.0 g O2 ×
2 mol C 3 H 3 N
53.06 g C 3 H 3 N
1 mol O 2
×
×
= 11.1 g C3H3N
32.00 g O 2
3 mol O 2
mol C 3 H 3 N
O2 produces the smallest amount of product; thus O2 is limiting, and 11.1 g C3H3N can be
produced.
123.
C2H6(g) + Cl2(g) → C2H5Cl(g) + HCl(g)
If C2H6 is limiting:
300. g C2H6 ×
1 mol C 2 H 6
1 mol C 2 H 5Cl 64.51 g C 2 H 5Cl
×
×
= 644 g C2H5Cl
30.07 g C 2 H 6
mol C 2 H 6
mol C 2 H 5Cl
If Cl2 is limiting:
650. g Cl2 ×
1 mol C 2 H 5Cl 64.51 g C 2 H 5Cl
1 mol Cl 2
×
×
= 591 g C2H5Cl
70.90 g Cl 2
mol Cl 2
mol C 2 H 5Cl
Cl2 is limiting because it produces the smaller quantity of product. Hence, the theoretical
yield for this reaction is 591 g C2H5Cl. The percent yield is:
percent yield =
124.
490. g
actual
× 100 =
× 100 = 82.9%
591 g
theoretical
a. 1142 g C6H5Cl ×
1 mol C 6 H 5Cl
1 mol C14 H 9 Cl5 354.46 g C14 H 9 Cl5
×
×
112.55 g C 6 H 5Cl
2 mol C 6 H 5Cl
mol C14 H 9 Cl5
= 1798 C14H9Cl5
74
CHAPTER 3
485 g C2HOCl3 ×
STOICHIOMETRY
1 mol C 2 HOCl3
1 mol C14 H 9 Cl5 354.46 g C14 H 9 Cl5
×
×
147.38 g C 2 HOCl3
mol C 2 HOCl3
mol C14 H 9 Cl5
= 1170 g C14H9Cl5
From the masses of product calculated, C2HOCl3 is limiting and 1170 g C14H9Cl5 can be
produced.
b. C2HOCl3 is limiting, and C6H5Cl is in excess.
c. 485 g C2HOCl3 ×
1 mol C 2 HOCl3
2 mol C 6 H 5Cl 112.55 g C 6 H 5Cl
×
×
147.38 g C 2 HOCl3
mol C 2 HOCl3
mol C 6 H 5Cl
= 741 g C6H5Cl reacted
1142 g − 741 g = 401 g C6H5Cl in excess
d. Percent yield =
125.
200.0 g DDT
× 100 = 17.1%
1170 g DDT
2.50 metric tons Cu3FeS3 ×
1 mol Cu 3 FeS3
3 mol Cu
1000 g
1000 kg
×
×
×
1 mol Cu 3 FeS3
metric ton
kg
342.71 g
×
1.39 × 106 g Cu (theoretical) ×
126.
63.55 g
= 1.39 × 106 g Cu (theoretical)
mol Cu
86.3 g Cu (actual)
= 1.20 × 106 g Cu = 1.20 × 103 kg Cu
100. g Cu ( theoretical)
= 1.20 metric tons Cu (actual)
P4(s) + 6 F2(g) → 4 PF3(g); the theoretical yield of PF3 is:
120. g PF3 (actual) ×
154 g PF3 ×
100.0 g PF3 ( theoretical)
= 154 g PF3 (theoretical)
78.1 g PF3 (actual)
1 mol PF3
6 mol F2
38.00 g F2
×
×
= 99.8 g F2
87.97 g PF3
4 mol PF3
mol F2
99.8 g F2 is needed to actually produce 120. g of PF3 if the percent yield is 78.1%.
Additional Exercises
127.
12
C21H6: 2(12.000000) + 6(1.007825) = 30.046950 u
12
C1H216O: 1(12.000000) + 2(1.007825) + 1(15.994915) = 30.010565 u
14
N16O: 1(14.003074) + 1(15.994915) = 29.997989 u
The peak results from 12C1H216O.
CHAPTER 3
128.
STOICHIOMETRY
75
We would see the peaks corresponding to:
10
B35Cl3 [mass ≈ 10 + 3(35) = 115 u], 10B35Cl237Cl (117), 10B35Cl37Cl2 (119),
10
B37Cl3 (121), 11B35Cl3 (116), 11B35Cl237Cl (118), 11B35Cl37Cl2 (120), 11B37Cl3 (122)
We would see a total of eight peaks at approximate masses of 115, 116, 117, 118, 119, 120,
121, and 122.
129.
0.368 g XeFn
Molar mass XeFn =
9.03 × 10
20
molecules XeFn ×
1 mol XeFn
= 245 g/mol
6.022 × 10 23 molecules
245 g = 131.3 g + n(19.00 g), n = 5.98; formula = XeF6
130.
a.
14 mol C ×
14.01 g
12.01 g
1.008 g
+ 18 mol H ×
+ 2 mol N ×
mol C
mol H
mol N
+ 5 mol O ×
1 mol C14 H18 N 2 O 5
= 3.40 × 10−2 mol C14H18N2O5
294.30 g C14 H18 N 2 O 5
b. 10.0 g C14H18N2O5 ×
c.
1.56 mol ×
d. 5.0 mg ×
16.00 g
= 294.30 g
mol O
294.3 g
= 459 g C14H18N2O5
mol
1 mol
6.022 × 10 23 molecules
1g
×
×
mol
1000 mg 294.30 g
= 1.0 × 1019 molecules C14H18N2O5
e. The chemical formula tells us that 1 molecule of C14H18N2O5 contains 2 atoms of N. If
we have 1 mole of C14H18N2O5 molecules, then 2 moles of N atoms are present.
1.2 g C14H18N2O5 ×
1 mol C14 H18 N 2 O 5
2 mol N
×
294.30 g C14 H18 N 2 O 5
mol C14 H18 N 2 O 5
×
f.
1.0 × 109 molecules ×
g. 1 molecule ×
131.
1 mol
6.022 × 10
1 mol
6.022 × 10
23
atoms
23
×
atoms
6.022 × 10 23 atoms N
= 4.9 × 1021 atoms N
mol N
×
294.30 g
= 4.9 × 10−13 g
mol
294.30 g
= 4.887 × 10−22 g C14H18N2O5
mol
Molar mass = 20(12.01) + 29(1.008) + 19.00 + 3(16.00) = 336.43 g/mol
Mass % C =
20(12.01) g C
× 100 = 71.40% C
336.43 g compound
76
CHAPTER 3
Mass % H =
Mass % F =
STOICHIOMETRY
29(1.008) g H
× 100 = 8.689% H
336.43 g compound
19.00 g F
× 100 = 5.648% F
336.43 g compound
Mass % O = 100.00 − (71.40 + 8.689 + 5.648) = 14.26% O or:
%O=
132.
3(16.00) g O
× 100 = 14.27% O
336.43 g compound
In 1 hour, the 1000. kg of wet cereal produced contains 580 kg H2O and 420 kg of cereal. We
want the final product to contain 20.% H2O. Let x = mass of H2O in final product.
x
= 0.20, x = 84 + (0.20)x, x = 105 ≈ 110 kg H2O
420 + x
The amount of water to be removed is 580 − 110 = 470 kg/h.
133.
Out of 100.00 g of adrenaline, there are:
56.79 g C ×
1 mol C
1 mol H
= 4.729 mol C; 6.56 g H ×
= 6.51 mol H
1.008 g H
12.01 g C
28.37 g O ×
1 mol O
1 mol N
= 1.773 mol O; 8.28 g N ×
= 0.591 mol N
16.00 g O
14.01 g N
Dividing each mole value by the smallest number:
4.729
6.51
1.773
0.591
= 8.00;
= 11.0;
= 3.00;
= 1.00
0.591
0.591
0.591
0.591
This gives adrenaline an empirical formula of C8H11O3N.
134.
Assuming 100.00 g of compound (mass hydrogen = 100.00 g − 49.31 g C − 43.79 g O
= 6.90 g H):
49.31 g C ×
1 mol C
12.01 g C
= 4.106 mol C; 6.90 g H ×
43.79 g O ×
1 mol O
16.00 g O
= 2.737 mol O
1 mol H
1.008 g H
= 6.85 mol H
Dividing all mole values by 2.737 gives:
2.737
4.106
6.85
= 1.500;
= 2.50;
= 1.000
2.737
2.737
2.737
Because a whole number ratio is required, the empirical formula is C3H5O2.
CHAPTER 3
STOICHIOMETRY
77
Empirical formula mass: 3(12.01) + 5(1.008) +2(16.00) = 73.07 g/mol
146.1
Molar mass
=
= 1.999; molecular formula = (C3H5O2)2 = C6H10O4
Empirical formula mass
73.07
135.
There are many valid methods to solve this problem. We will assume 100.00 g of compound,
and then determine from the information in the problem how many moles of compound
equals 100.00 g of compound. From this information, we can determine the mass of one
mole of compound (the molar mass) by setting up a ratio. Assuming 100.00 g cyanocobalamin:
1 mol Co
1 mol cyanocobalamin
×
mol cyanocobalamin = 4.34 g Co ×
58.93 g Co
mol Co
= 7.36 × 10 −2 mol cyanocobalamin
x g cyanocobalamin
100.00 g
=
, x = molar mass = 1.36 × 103 g/mol
−2
1 mol cyanocobalamin
7.36 × 10 mol
136.
2 tablets ×
0.262 g C 7 H 5 BiO 4
1 mol C 7 H 5 BiO 4
1 mol Bi
209.0 g Bi
×
×
×
tablet
362.11 g C 7 H 5 BiO 4 1 mol C 7 H 5 BiO 4
mol Bi
= 0.302 g Bi consumed
137.
Empirical formula mass = 12.01 + 1.008 = 13.02 g/mol; because 104.14/13.02 = 7.998 ≈ 8,
the molecular formula for styrene is (CH)8 = C8H8.
2.00 g C8H8 ×
138.
1 mol C 8 H 8
8 mol H
6.022 × 10 23 atoms H
= 9.25 × 1022 atoms H
×
×
mol C 8 H 8
mol H
104.14 g C 8 H 8
41.98 mg CO2 ×
6.45 mg H2O ×
12.01 mg C
11.46 mg
= 11.46 mg C; % C =
× 100 = 57.85% C
44.01 mg CO 2
19.81 mg
2.01 6 mg H
0.772 mg
= 0.722 mg H; % H =
× 100 = 3.64% H
19.81 mg
18.02 mg H 2 O
% O = 100.00 − (57.85 + 3.64) = 38.51% O
Out of 100.00 g terephthalic acid, there are:
57.85 g C ×
1 mol H
1 mol C
= 4.817 mol C; 3.64 g H ×
= 3.61 mol H
1.008 g H
12.01 g C
38.51 g O ×
1 mol O
= 2.407 mol O
16.00 g O
4.817
3.61
= 2.001;
= 1.50;
2.407
2.407
2.407
= 1.000
2.407
78
CHAPTER 3
STOICHIOMETRY
The C : H : O mole ratio is 2 : 1.5 : 1 or 4 : 3 : 2. The empirical formula is C4H3O2.
Mass of C4H3O2 ≈ 4(12) + 3(1) + 2(16) = 83g/mol.
Molar mass =
139.
17.3 g H ×
41.5 g
166
= 166 g/mol;
= 2.0; the molecular formula is C8H6O4.
0.250 mol
83
1 mol C
1 mol H
= 17.2 mol H; 82.7 g C ×
= 6.89 mol C
1.008 g H
12.01 g C
17.2
= 2.50; the empirical formula is C2H5.
6.89
The empirical formula mass is ~29 g/mol, so two times the empirical formula would put the
compound in the correct range of the molar mass. Molecular formula = (C2H5)2 = C4H10.
2.59 × 1023 atoms H ×
1 molecule C 4 H10
1 mol C 4 H10
= 4.30 × 10 −2 mol C4H10
×
23
10 atoms H
6.022 × 10 molecules
4.30 × 10 −2 mol C4H10 ×
140.
58.12 g
= 2.50 g C4H10
mol C 4 H 10
Assuming 100.00 g E3H8:
mol E = 8.73 g H ×
1 mol H
3 mol E
×
= 3.25 mol E
1.008 g H 8 mol H
xgE
91.27 g E
=
, x = molar mass of E = 28.1 g/mol; atomic mass of E = 28.1 u
1 mol E
3.25 mol E
Note: From the periodic table, element E is silicon, Si.
141.
Mass of H2O = 0.755 g CuSO4•xH2O − 0.483 g CuSO4 = 0.272 g H2O
0.483 g CuSO4 ×
0.272 g H2O ×
1 mol CuSO 4
= 0.00303 mol CuSO4
159.62 g CuSO 4
1 mol H 2 O
= 0.0151 mol H2O
18.02 g H 2 O
4.98 mol H 2 O
0.0151 mol H 2 O
=
; compound formula = CuSO4•5H2O, x = 5
0.00303 g CuSO 4
1 mol CuSO 4
142.
a. Only acrylonitrile contains nitrogen. If we have 100.00 g of polymer:
8.80 g N ×
1 mol C 3 H 3 N 53.06 g C 3 H 3 N
×
= 33.3 g C3H3N
14.01 g N
1 mol C 3 H 3 N
CHAPTER 3
STOICHIOMETRY
% C3H3N =
79
33.3 g C 3 H 3 N
= 33.3% C3H3N
100.00 g polymer
Only butadiene in the polymer reacts with Br2:
0.605 g Br2 ×
% C4H6 =
1 mol C 4 H 6
54.09 g C 4 H 6
1 mol Br2
×
×
= 0.205 g C4H6
159.8 g Br2
mol Br2
mol C 4 H 6
0.205 g
× 100 = 17.1% C4H6
1.20 g
b. If we have 100.0 g of polymer:
33.3 g C3H3N ×
1 mol C 3 H 3 N
= 0.628 mol C3H3N
53.06 g
17.1 g C4H6 ×
1 mol C 4 H 6
= 0.316 mol C4H6
54.09 g C 4 H 6
49.6 g C8H8 ×
1 mol C8 H 8
= 0.476 mol C8H8
104.14 g C8 H 8
Dividing by 0.316:
0.316
0.476
0.628
= 1.99;
= 1.00;
= 1.51
0.316
0.316
0.316
This is close to a mole ratio of 4 : 2 : 3. Thus there are 4 acrylonitrile to 2 butadiene to 3
styrene molecules in the polymer, or (A4B2S3)n.
143.
1.20 g CO2 ×
1 mol C 24 H 30 N 3O
1 mol CO 2
1 mol C
376.51 g
×
×
×
44.01 g
mol CO 2
24 mol C
mol C 24 H 30 N 3O
= 0.428 g C24H30N3O
0.428 g C 24 H 30 N 3 O
× 100 = 42.8% C24H30N3O (LSD)
1.00 g sample
144.
a. CH4(g) + 4 S(s) → CS2(l) + 2 H2S(g) or 2 CH4(g) + S8(s) → 2 CS2(l) + 4 H2S(g)
b. 120. g CH4 ×
120. g S ×
1 mol CH 4
1 mol CS 2
76.15 g CS 2
×
×
= 570. g CS2
16.04 g CH 4
mol CH 4
mol CS 2
1 mol CS 2
76.15 g CS 2
1 mol S
×
×
= 71.2 g CS2
32.07 g S
4 mol S
mol CS 2
Because S produces the smaller quantity of CS2, sulfur is the limiting reactant and 71.2 g
CS2 can be produced. The same amount of CS2 would be produced using the balanced
equation with S8.
80
145.
CHAPTER 3
126 g B5H9 ×
192 g O2 ×
STOICHIOMETRY
1 mol B5 H 9
9 mol H 2 O 18.02 g H 2 O
×
×
= 162 g H2O
63.12 g B5 H 9
2 mol B5 H 9
mol H 2 O
1 mol O 2
9 mol H 2 O 18.02 g H 2 O
×
×
= 81.1 g H2O
32.00 g O 2
12 mol O 2
mol H 2 O
Because O2 produces the smallest quantity of product, O2 is limiting and 81.1g H2O can be
produced.
146.
2 NaNO3(s) → 2 NaNO2(s) + O2(g); the amount of NaNO3 in the impure sample is:
0.2864 g NaNO2 ×
2 mol NaNO3
85.00 g NaNO3
1 mol NaNO 2
×
×
69.00 g NaNO 2
2 mol NaNO 2
mol NaNO3
= 0.3528 g NaNO3
0.3528 g NaNO3
Mass percent NaNO3 =
× 100 = 83.40%
0.4230 g sample
147.
453 g Fe ×
1 mol Fe 2 O 3 159.70 g Fe 2 O 3
1 mol Fe
×
×
= 648 g Fe2O3
mol Fe 2 O 3
2 mol Fe
55.85 g Fe
Mass percent Fe2O3 =
148.
648 g Fe 2 O 3
× 100 = 86.2%
752 g ore
a. Mass of Zn in alloy = 0.0985 g ZnCl2 ×
% Zn =
65.38 g Zn
= 0.0473 g Zn
136.28 g ZnCl 2
0.0473 g Zn
× 100 = 9.34% Zn; % Cu = 100.00 − 9.34 = 90.66% Cu
0.5065 g brass
b. The Cu remains unreacted. After filtering, washing, and drying, the mass of the unreacted
copper could be measured.
149.
Assuming 1 mole of vitamin A (286.4 g vitamin A):
mol C = 286.4 g vitamin A ×
mol H = 286.4 g vitamin A ×
0.8386 g C
1 mol C
×
= 20.00 mol C
g vitamin A
12.01 g C
0.1056 g H
1 mol H
×
= 30.00 mol H
g vitamin A
1.008 g H
Because 1 mole of vitamin A contains 20 mol C and 30 mol H, the molecular formula of
vitamin A is C20H30E. To determine E, let’s calculate the molar mass of E:
286.4 g = 20(12.01) + 30(1.008) + molar mass E, molar mass E = 16.0 g/mol
From the periodic table, E = oxygen, and the molecular formula of vitamin A is C20H30O.
CHAPTER 3
150.
STOICHIOMETRY
81
a. At 40.0 g of Na added, Cl2 and Na both run out at the same time (both are limiting reactants). Past 40.0 g of Na added, Cl2 is limiting, and because the amount of Cl2 present in
each experiment was the same quantity, no more NaCl can be produced. Before 40.0 g of
Na added, Na was limiting. As more Na was added (up to 40.0 g Na), more NaCl was
produced.
b. 20.0 g Na ×
1 mol Na
2 mol NaCl 58.44 g NaCl
×
×
= 50.8 g NaCl
22.99 g Na
2 mol Na
mol NaCl
c. At 40.0 g Na added, both Cl2 and Na are present in stoichiometric amounts.
40.0 g Na ×
1 mol Cl 2
70.90 g Cl 2
1 mol Na
×
×
= 61.7 g Cl2
22.99 g Na 2 mol Na
mol Cl 2
61.7 g Cl2 was present at 40.0 g Na added, and from the problem, the same 61.7 g Cl2 was
present in each experiment.
d. At 50.0 g Na added, Cl2 is limiting:
61.7 g Cl2 ×
e. 20.0 g Na ×
1 mol Cl 2
2 mol NaCl 58.44 g NaCl
×
×
= 101.7 g = 102 g NaCl
70.90 g Cl 2
1 mol Cl 2
mol NaCl
1 mol Cl 2
70.90 g Cl 2
1 mol Na
×
×
= 30.8 g Cl2 reacted
22.99 g Na 2 mol Na
mol Cl 2
Excess Cl2 = 61.7 g Cl2 initially – 30.8 g Cl2 reacted = 30.9 g Cl2 in excess
Note: We know that 40.0 g Na is the point where Na and the 61.7 g of Cl2 run out at the
same time. So if 20.0 g of Na are reacted, one-half of the Cl2 that was present at 40.0 g
Na reacted will be in excess. The previous calculation confirms this.
For 50.0 g Na reacted, Cl2 is limiting and 40.0 g Na will react as determined previously.
Excess Na = 50.0 g Na initially – 40.0 g Na reacted = 10.0 g Na in excess.
151.
X2Z: 40.0% X and 60.0% Z by mass;
40.0/A x
(40.0)A z
mol X
=2=
=
or Az = 3Ax,
mol Z
60.0/A z
(60.0)A x
where A = molar mass.
For XZ2, molar mass = Ax + 2Az = Ax + 2(3Ax) = 7Ax.
Mass percent X =
Ax
× 100 = 14.3% X; % Z = 100.0 − 14.3 = 85.7% Z
7A x
ChemWork Problems
The answers to the problems 152-159 (or a variation to these problems) are found in OWL. These
problems are also assignable in OWL.
82
CHAPTER 3
STOICHIOMETRY
Challenge Problems
160.
GaAs can be either 69GaAs or 71GaAs. The mass spectrum for GaAs will have 2 peaks at 144
(= 69 + 75) and 146 (= 71 + 75) with intensities in the ratio of 60 : 40 or 3 : 2.
144
146
Ga2As2 can be 69Ga2As2, 69Ga71GaAs2, or 71Ga2As2. The mass spectrum will have 3 peaks at
288, 290, and 292 with intensities in the ratio of 36 : 48 : 16 or 9 : 12 : 4. We get this ratio
from the following probability table:
69
Ga (0.60)
Ga (0.40)
69
0.36
0.24
71
0.24
0.16
Ga (0.60)
Ga (0.40)
288
161.
71
290
292
The volume of a gas is proportional to the number of molecules of gas. Thus the formulas
are:
I: NH3 ;
II: N2H4;
III: HN3
The mass ratios are:
I:
4.634 g N
82.25 g N
=
;
17.75 g H
gH
II:
6.949 g N
;
gH
III:
41.7 g N
gH
If we set the atomic mass of H equal to 1.008, then the atomic mass, A, for nitrogen is:
I: 14.01;
II: 14.01;
For example, for compound I:
85
162.
87
Rb atoms
= 2.591;
Rb atoms
III. 14.0
A
4.634
=
, A = 14.01
3(1.008)
1
if we had exactly 100 atoms, x = number of 85Rb atoms, and
100 − x = number of 87Rb atoms.
CHAPTER 3
STOICHIOMETRY
83
x
259.1
= 2.591, x = 259.1 − (2.591)x, x =
= 72.15; 72.15% 85Rb
100 − x
3.591
0.7215(84.9117) + 0.2785(A) = 85.4678, A =
163.
85.4678 − 61.26
= 86.92 u
0.2785
First, we will determine composition in mass percent. We assume that all the carbon in the
0.213 g CO2 came from the 0.157 g of the compound and that all the hydrogen in the 0.0310
g H2O came from the 0.157 g of the compound.
0.213 g CO2 ×
0.0581 g C
12.01 g C
= 0.0581 g C; % C =
× 100 = 37.0% C
44.01 g CO 2
0.157 g compound
0.0310 g H2O ×
2.016 g H
3.47 × 10 −3 g
= 3.47 × 10−3 g H; % H =
× 100 = 2.21% H
18.02 g H 2 O
0.157 g
We get the mass percent of N from the second experiment:
0.0230 g NH3 Η
%N=
14.01 g N
= 1.89 × 10−2 g N
17.03 g NH 3
1.89 × 10 −2 g
× 100 = 18.3% N
0.103 g
The mass percent of oxygen is obtained by difference:
% O = 100.00 − (37.0 + 2.21 + 18.3) = 42.5% O
So, out of 100.00 g of compound, there are:
37.0 g C ×
1 mol C
1 mol H
= 3.08 mol C; 2.21 g H ×
= 2.19 mol H
1.008 g H
12.01 g C
18.3 g N ×
1 mol N
1 mol O
= 1.31 mol N; 42.5 g O ×
= 2.66 mol O
16.00 g O
14.01 g N
Lastly, and often the hardest part, we need to find simple whole-number ratios. Divide all
mole values by the smallest number:
2.66
2.19
3.08
1.31
= 2.35;
= 1.67;
= 1.00;
= 2.03
1.31
1.31
1.31
1.31
Multiplying all these ratios by 3 gives an empirical formula of C7H5N3O6.
84
164.
CHAPTER 3
1.0 × 106 kg HNO3 ×
STOICHIOMETRY
1000 g HNO3
1 mol HNO3
×
= 1.6 × 107 mol HNO3
kg HNO3
63.02 g HNO3
We need to get the relationship between moles of HNO3 and moles of NH3. We have to use
all three equations.
16 mol HNO3
2 mol HNO3
2 mol NO 2
4 mol NO
×
×
=
3 mol NO 2
2 mol NO
4 mol NH 3
24 mol NH 3
Thus we can produce 16 mol HNO3 for every 24 mol NH3, we begin with:
1.6 × 107 mol HNO3 ×
24 mol NH 3
17.03 g NH 3
×
= 4.1 × 108 g or 4.1 × 105 kg NH3
16 mol HNO3
mol NH 3
This is an oversimplified answer. In practice, the NO produced in the third step is recycled
back continuously into the process in the second step. If this is taken into consideration, then
the conversion factor between mol NH3 and mol HNO3 turns out to be 1 : 1; that is, 1 mole of
NH3 produces 1 mole of HNO3. Taking into consideration that NO is recycled back gives an
answer of 2.7 × 105 kg NH3 reacted.
165.
Fe(s) +
1
2
O 2 (g ) → FeO(s) ; 2 Fe(s) +
20.00 g Fe ×
3
2
O 2 (g ) → Fe 2 O 3 (s)
1 mol Fe
= 0.3581 mol
55.85 g
(11.20 − 3.24) g O 2 ×
1 mol O 2
= 0.2488 mol O2 consumed (1 extra sig. fig.)
32.00 g
Let’s assume x moles of Fe reacts to form x moles of FeO. Then 0.3581 – x, the remaining
moles of Fe, reacts to form Fe2O3. Balancing the two equations in terms of x:
1
x O 2 → x FeO
2
3  0.3581 − x 
 0.3581 − x 
(0.3581 − x) mol Fe + 
 mol Fe 2 O 3
 mol O 2 → 
2 
2
2



x Fe +
Setting up an equation for total moles of O2 consumed:
1
2
x +
3
4
(0.3581 − x) = 0.2488 mol O 2 , x = 0.0791 = 0.079 mol FeO
0.079 mol FeO ×
71.85 g FeO
= 5.7 g FeO produced
mol
Mol Fe2O3 produced =
0.140 mol Fe2O3 ×
0.3581 − 0.079
= 0.140 mol Fe2O3
2
159.70 g Fe 2 O 3
= 22.4 g Fe2O3 produced
mol
CHAPTER 3
166.
STOICHIOMETRY
85
2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(l); C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(l)
30.07 g/mol
44.09 g/mol
Let x = mass C2H6, so 9.780 − x = mass C3H8. Use the balanced equation to set up a
mathematical expression for the moles of O2 required.
x
7
9.780 − x 5
×
+
× = 1.120 mol O2
30.07 2
44.09
1
Solving: x = 3.7 g C2H6;
167.
3.7 g
× 100 = 38% C2H6 by mass
9.780 g
The two relevant equations are:
Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g) and Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
Let x = mass Mg, so 10.00 − x = mass Zn. From the balanced equations, moles H2 produced =
moles Zn reacted + moles Mg reacted.
Mol H2 = 0.5171 g H2 ×
0.2565 =
1 mol H 2
= 0.2565 mol H2
2.016 g H 2
10.00 − x
x
+
; solving: x = 4.008 g Mg
24.31
65.38
4.008 g
× 100 = 40.08% Mg
10.00 g
168.
a N2H4 + b NH3 + (10.00 − 4.062) O2 → c NO2 + d H2O
Setting up four equations to solve for the four unknowns:
2a + b = c
(N mol balance)
2c + d = 2(10.00 − 4.062)
(O mol balance)
4a + 3b = 2d
(H mol balance)
a(32.05) + b(17.03) = 61.00
(mass balance)
Solving the simultaneous equations gives a = 1.12 = 1.1 mol N2H4.
1.1 mol N 2 H 4 × 32.05 g / mol N 2 H 4
× 100 = 58% N2H4
61.00 g
169.
We know that water is a product, so one of the elements in the compound is hydrogen.
XaHb + O2 → H2O + ?
86
CHAPTER 3
To balance the H atoms, the mole ratio between XaHb and H2O =
Mol compound =
STOICHIOMETRY
2
.
b
1.21 g
1.39 g
= 0.0224 mol; mol H2O =
= 0.0671 mol
62.09 g / mol
18.02 g / mol
2
0.0224
, b = 6; XaH6 has a molar mass of 62.09 g/mol.
=
b
0.0671
62.09 = a(molar mass of X) + 6(1.008), a(molar mass of X) = 56.04
Some possible identities for X could be Fe (a = 1), Si (a = 2), N (a = 4), and Li (a = 8). N fits
the data best, so N4H6 is the most likely formula.
170.
The balanced equation is 2 Sc(s) + 2x HCl(aq) → 2 ScClx(aq)+ x H2(g)
The mole ratio of Sc : H2 =
Mol Sc = 2.25 g Sc ×
2
.
x
1 mol Sc
= 0.0500 mol Sc
44.96 g Sc
Mol H2 = 0.1502 g H2 ×
1 mol H 2
= 0.07450 mol H2
2.016 g H 2
0.0500
2
=
, x = 3; the formula is ScCl3.
x
0.07450
171.
Total mass of copper used:
10,000 boards ×
(8.0 cm × 16.0 cm × 0.060 cm) 8.96 g
×
= 6.9 × 105 g Cu
3
board
cm
Amount of Cu to be recovered = 0.80 × (6.9 × 105 g) = 5.5 × 105 g Cu.
5.5 × 105 g Cu ×
1 mol Cu(NH 3 ) 4 Cl 2
202.59 g Cu(NH 3 ) 4 Cl 2
1 mol Cu
×
×
63.55 g Cu
mol Cu
mol Cu(NH 3 ) 4 Cl 2
= 1.8 × 106 g Cu(NH3)4Cl2
5.5 × 105 g Cu ×
172.
4 mol NH 3 17.03 g NH 3
1 mol Cu
×
×
= 5.9 × 105 g NH3
mol NH 3
63.55 g Cu
mol Cu
a. From the reaction stoichiometry we would expect to produce 4 mol of acetaminophen for
every 4 mol of C6H5O3N reacted. The actual yield is 3 mol of acetaminophen compared
to a theoretical yield of 4 mol of acetaminophen. Solving for percent yield by mass
(where M = molar mass acetaminophen):
percent yield =
3 mol × M
× 100 = 75%
4 mol × M
CHAPTER 3
STOICHIOMETRY
87
b. The product of the percent yields of the individual steps must equal the overall yield,
75%.
(0.87)(0.98)(x) = 0.75, x = 0.88; step III has a percent yield of 88%.
173.
10.00 g XCl2 + excess Cl2 → 12.55 g XCl4; 2.55 g Cl reacted with XCl2 to form XCl4. XCl4
contains 2.55 g Cl and 10.00 g XCl2. From the mole ratios, 10.00 g XCl2 must also contain
2.55 g Cl; mass X in XCl2 = 10.00 − 2.55 = 7.45 g X.
2.55 g Cl ×
1 mol XCl 2
1 mol Cl
1 mol X
×
×
= 3.60 × 10 −2 mol X
35.45 g Cl
2 mol Cl
mol XCl 2
So 3.60 × 10 −2 mol X has a mass equal to 7.45 g X. The molar mass of X is:
7.45 g X
= 207 g/mol X; atomic mass = 207 u, so X is Pb.
3.60 × 10 − 2 mol X
174.
4.000 g M2S3 → 3.723 g MO2
There must be twice as many moles of MO2 as moles of M2S3 in order to balance M in the
reaction. Setting up an equation for 2(mol M2S3) = mol MO2 where A = molar mass M:


4.000 g
3.723 g
8.000
3.723
=
2
,
 =
A + 32.00
 2A + 3(32.07)  A + 2(16.00) 2A + 96.21
(8.000)A + 256.0 = (7.446)A + 358.2, (0.554)A = 102.2, A = 184 g/mol; atomic mass
= 184 u
Note: From the periodic table, M is tungsten, W.
175.
Consider the case of aluminum plus oxygen. Aluminum forms Al3+ ions; oxygen forms O2−
anions. The simplest compound between the two elements is Al2O3. Similarly, we would
expect the formula of any Group 6A element with Al to be Al2X3. Assuming this, out of
100.00 g of compound, there are 18.56 g Al and 81.44 g of the unknown element, X. Let’s
use this information to determine the molar mass of X, which will allow us to identify X from
the periodic table.
18.56 g Al ×
1 mol Al
3 mol X
×
= 1.032 mol X
26.98 g Al 2 mol Al
81.44 g of X must contain 1.032 mol of X.
Molar mass of X =
81.44 g X
= 78.91 g/mol X.
1.032 mol X
From the periodic table, the unknown element is selenium, and the formula is Al2Se3.
88
176.
CHAPTER 3
STOICHIOMETRY
Let x = mass KCl and y = mass KNO3. Assuming 100.0 g of mixture, x + y = 100.0 g.
Molar mass KCl = 74.55 g/mol; molar mass KNO3 = 101.11 g/mol
Mol KCl =
x
y
; mol KNO3 =
74.55
101.11
Knowing that the mixture is 43.2% K, then in the 100.0 g mixture, an expression for the mass
of K is:
y 
 x
39.10 
+
 = 43.2
 74.55 101.11 
We have two equations and two unknowns:
(0.5245)x + (0.3867)y = 43.2
x +
y = 100.0
Solving, x = 32.9 g KCl;
177.
32.9 g
× 100 = 32.9% KCl
100.0 g
The balanced equations are:
4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g) and 4 NH3(g) + 7 O2(g) → 4 NO2(g)
+ 6 H2O(g)
Let 4x = number of moles of NO formed, and let 4y = number of moles of NO2 formed.
Then:
4x NH3 + 5x O2 → 4x NO + 6x H2O and 4y NH3 + 7y O2 → 4y NO2 + 6y H2O
All the NH3 reacted, so 4x + 4y = 2.00. 10.00 − 6.75 = 3.25 mol O2 reacted, so 5x + 7y
= 3.25.
Solving by the method of simultaneous equations:
20x + 28y = 13.0
−20x − 20y = −10.0
8y = 3.0, y = 0.38; 4x + 4 × 0.38 = 2.00, x = 0.12
Mol NO = 4x = 4 × 0.12 = 0.48 mol NO formed
178.
CxHyOz + oxygen → x CO2 + y/2 H2O
2.20 g CO 2 ×
Mass % C in aspirin =
1 mol C
1 mol CO 2
12.01 g C
×
×
44.01 g CO 2
mol CO 2
mol C
= 60.0% C
1.00 g aspirin
CHAPTER 3
STOICHIOMETRY
89
0.400 g H 2 O ×
Mass % H in aspirin =
1 mol H 2 O
2 mol H
1.008 g H
×
×
18.02 g H 2 O mol H 2 O
mol H
= 4.48% H
1.00 g aspirin
Mass % O = 100.00 − (60.0 + 4.48) = 35.5% O
Assuming 100.00 g aspirin:
60.0 g C ×
1 mol C
1 mol H
= 5.00 mol C; 4.48 g H ×
= 4.44 mol H
12.01 g C
1.008 g H
35.5 g O ×
1 mol O
= 2.22 mol O
16.00 g O
Dividing by the smallest number:
5.00
4.44
= 2.25;
= 2.00
2.22
2.22
Empirical formula: (C2.25 H2.00O)4 = C9H8O4. Empirical mass ≈ 9(12) + 8(1) + 4(16)
= 180 g/mol; this is in the 170–190 g/mol range, so the molecular formula is also C9H8O4.
Balance the aspirin synthesis reaction to determine the formula for salicylic acid.
CaHbOc + C4H6O3 → C9H8O4 + C2H4O2, CaHbOc = salicylic acid = C7H6O3
Integrative Problems
179.
1 mol Fe
6.022 × 10 23 atoms Fe
= 113 atoms Fe
×
55.85 g Fe
mol Fe
a. 1.05 × 10 −20 g Fe ×
b. The total number of platinum atoms is 14 × 20 = 280 atoms (exact number). The mass of
these atoms is:
280 atoms Pt ×
1 mol Pt
6.022 × 10
c. 9.071 × 10 −20 g Ru ×
180.
23
atoms Pt
×
195.1 g Pt
= 9.071 × 10 −20 g Pt
mol Pt
1 mol Ru
6.022 × 10 23 atoms Ru
= 540.3 = 540 atoms Ru
×
101.1 g Ru
mol Ru
Assuming 100.00 g of tetrodotoxin:
41.38 g C ×
5.37 g H ×
1 mol C
1 mol N
= 3.445 mol C; 13.16 g N ×
= 0.9393 mol N
12.01 g C
14.01 g N
1 mol H
1 mol O
= 5.33 mol H; 40.09 g O ×
= 2.506 mol O
1.008 g H
16.00 g O
90
CHAPTER 3
STOICHIOMETRY
Divide by the smallest number:
3.445
5.33
2.506
= 3.668;
= 5.67;
= 2.668
0.9393
0.9393
0.9393
To get whole numbers for each element, multiply through by 3.
Empirical formula: (C3.668H5.67NO2.668)3 = C11H17N3O8; the mass of the empirical formula is
319.3 g/mol.
Molar mass tetrodotoxin =
1.59 × 10 −21 g
= 319 g/mol
1 mol
3 molecules ×
6.022 × 10 23 molecules
Because the empirical mass and molar mass are the same, the molecular formula is the same
as the empirical formula, C11H17N3O8.
165 lb ×
1 kg
1 × 10 −6 g
1 mol
10. μg
6.022 × 10 23 molecules
×
×
×
×
2.2046 lb
kg
μg
319.3 g
1 mol
= 1.4 × 1018 molecules tetrodotoxin is the LD50 dosage
181.
0.105 g
Molar mass X2 =
8.92 × 10
20
1 mol
molecules ×
6.022 × 10 23 molecules
= 70.9 g/mol
The mass of X = 1/2(70.9 g/mol) = 35.5 g/mol. This is the element chlorine.
Assuming 100.00 g of MX3 (= MCl3) compound:
54.47 g Cl ×
1 mol
= 1.537 mol Cl
35.45 g
1.537 mol Cl ×
Molar mass of M =
1 mol M
= 0.5123 mol M
3 mol Cl
45.53 g M
= 88.87 g/mol M
0.5123 mol M
M is the element yttrium (Y), and the name of YCl3 is yttrium(III) chloride.
The balanced equation is 2 Y + 3 Cl2 → 2 YCl3.
Assuming Cl2 is limiting:
1.00 g Cl2 ×
2 mol YCl3
195.26 g YCl3
1 mol Cl 2
×
×
= 1.84 g YCl3
70.90 g Cl 2
3 mol Cl 2
1 mol YCl3
CHAPTER 3
STOICHIOMETRY
91
Assuming Y is limiting:
2 mol YCl3
195.26 g YCl3
1 mol Y
×
×
= 2.20 g YCl3
88.91 g Y
2 mol Y
1 mol YCl3
1.00 g Y ×
Because Cl2, when it all reacts, produces the smaller amount of product, Cl2 is the limiting
reagent, and the theoretical yield is 1.84 g YCl3.
182.
2 As + 4 AsI3 → 3 As2I4
Volume of As cube = (3.00 cm)3 = 27.0 cm3
27.0 cm3 ×
5.72 g As
cm
3
×
3 mol As 2 I 4
657.44 g As 2 I 4
1 mol As
×
×
= 2030 g As2I4
74.92 g As
2 mol As
mol As 2 I 4
1.01 × 1024 molecules AsI3 ×
1 mol AsI 3
6.022 × 10
23
molecules AsI 3
×
×
3 mol As 2 I 4
4 mol AsI 3
657.44 g As 2 I 4
= 827 g As 2 I 4
mol As 2 I 4
Because the reactant AsI3 produces the smaller quantity of product, then AsI3 is the limiting
reactant and 827 g As2I4 is the theoretical yield.
0.756 =
actual yield
, actual yield = 0.756 × 827 g = 625 g As2I4
827 g
Marathon Problems
183.
Let M = unknown element; mass O in oxide = 3.708 g – 2.077 g = 1.631 g O
In 3.708 g of compound:
1.631 g O ×
1 mol O
= 0.1019 g mol O
16.00 g O
If MO is the formula of the oxide, then M has a molar mass of
2.077 g M
= 20.38 g/mol.
0.1019 mol M
This is too low for the molar mass. We must have fewer moles of M than moles O present in
the formula. Some possibilities are MO2, M2O3, MO3, etc. It is a guessing game as to which
to try. Let’s assume an MO2 formula. Then the molar mass of M is:
2.077 g M
= 40.77 g/mol
1 mol M
0.1019 mol O ×
2 mol O
This is close to calcium, but calcium forms an oxide having the CaO formula, not CaO2.
92
CHAPTER 3
STOICHIOMETRY
If MO3 is assumed to be the formula, then the molar mass of M calculates to be 61.10 g/mol,
which is too large. Therefore, the mol O to mol M ratio must be between 2 and 3. Some
reasonable possibilities are 2.25, 2.33, 2.5, 2.67, and 2.75 (these are reasonable because they
will lead to whole number formulas). Trying a mol O to mol M ratio of 2.5 : 1 gives a molar
mass of:
2.077 g M
= 50.96 g/mol
1 mol M
0.1019 mol O ×
2.5 mol O
This is the molar mass of vanadium, and V2O5 is a reasonable formula for an oxide of
vanadium. The other choices for the O : M mole ratios between 2 and 3 do not give as
reasonable results. Therefore, M is most likely vanadium, and the formula is V2O5.
184.
a. i.
If the molar mass of A is greater than the molar mass of B, then we cannot determine
the limiting reactant because, while we have a fewer number of moles of A, we also
need fewer moles of A (from the balanced reaction).
ii. If the molar mass of B is greater than the molar mass of A, then B is the limiting
reactant because we have a fewer number of moles of B and we need more B (from
the balanced reaction).
b. A + 5 B → 3 CO2 + 4 H2O
To conserve mass: 44.01 + 5(B) = 3(44.01) + 4(18.02); solving: B = 32.0 g/mol
Because B is diatomic, the best choice for B is O2.
c. We can solve this without mass percent data simply by balancing the equation:
A + 5 O2 → 3 CO2 + 4 H2O
A must be C3H8 (which has a similar molar mass to CO2). This is also the empirical
formula.
Note:
3(12.01)
× 100 = 81.71% C. So this checks.
3(12.01) + 8(1.008)
CHAPTER 4
TYPES OF CHEMICAL REACTIONS AND SOLUTION
STOICHIOMETRY
Questions
13.
a. Polarity is a term applied to covalent compounds. Polar covalent compounds have an
unequal sharing of electrons in bonds that results in unequal charge distribution in the
overall molecule. Polar molecules have a partial negative end and a partial positive end.
These are not full charges as in ionic compounds but are charges much smaller in
magnitude. Water is a polar molecule and dissolves other polar solutes readily. The
oxygen end of water (the partial negative end of the polar water molecule) aligns with the
partial positive end of the polar solute, whereas the hydrogens of water (the partial
positive end of the polar water molecule) align with the partial negative end of the solute.
These opposite charge attractions stabilize polar solutes in water. This process is called
hydration. Nonpolar solutes do not have permanent partial negative and partial positive
ends; nonpolar solutes are not stabilized in water and do not dissolve.
b. KF is a soluble ionic compound, so it is a strong electrolyte. KF(aq) actually exists as
separate hydrated K+ ions and hydrated F− ions in solution: C6H12O6 is a polar covalent
molecule that is a nonelectrolyte. C6H12O6 is hydrated as described in part a.
c. RbCl is a soluble ionic compound, so it exists as separate hydrated Rb+ ions and hydrated
Cl− ions in solution. AgCl is an insoluble ionic compound, so the ions stay together in
solution and fall to the bottom of the container as a precipitate.
d. HNO3 is a strong acid and exists as separate hydrated H+ ions and hydrated NO3− ions in
solution. CO is a polar covalent molecule and is hydrated as explained in part a.
14.
2.0 L × 3.0 mol/L = 6.0 mol HCl; the 2.0 L of solution contains 6.0 mol of the solute. HCl
is a strong acid; it exists in aqueous solution as separate hydrated H+ ions and hydrated Cl−
ions. So the solution will contain 6.0 mol of H+(aq) and 6.0 mol of Cl− (aq). For the acetic
acid solution, HC2H3O2 is a weak acid instead of a strong acid. Only some of the 6.0 moles
of HC2H3O2 molecules will dissociate into H+(aq) and C2H3O2−(aq). The 2.0 L of 3.0 M
HC2H3O2 solution will contain mostly hydrated HC2H3O2 molecules but will also contain
some hydrated H+ ions and hydrated C2H3O2− ions.
15.
Only statement b is true. A concentrated solution can also contain a nonelectrolyte dissolved
in water, e.g., concentrated sugar water. Acids are either strong or weak electrolytes. Some
ionic compounds are not soluble in water, so they are not labeled as a specific type of
electrolyte.
93
94
16.
CHAPTER 4
SOLUTION STOICHIOMETRY
One mole of NaOH dissolved in 1.00 L of solution will produce 1.00 M NaOH. First, weigh
out 40.00 g of NaOH (1.000 mol). Next, add some water to a 1-L volumetric flask (an
instrument that is precise to 1.000 L). Dissolve the NaOH in the flask, add some more water,
mix, add more water, mix, etc. until water has been added to 1.000-L mark of the volumetric
flask. The result is 1.000 L of a 1.000 M NaOH solution. Because we know the volume to
four significant figures as well as the mass, the molarity will be known to four significant
figures. This is good practice, if you need a three-significant-figure molarity, your
measurements should be taken to four significant figures.
When you need to dilute a more concentrated solution with water to prepare a solution, again
make all measurements to four significant figures to ensure three significant figures in the
molarity. Here, we need to cut the molarity in half from 2.00 M to 1.00 M. We would start
with 1 mole of NaOH from the concentrated solution. This would be 500.0 mL of 2.00 M
NaOH. Add this to a 1-L volumetric flask with addition of more water and mixing until the
1.000-L mark is reached. The resulting solution would be 1.00 M.
17.
Use the solubility rules in Table 4.1. Some soluble bromides by Rule 2 would be NaBr, KBr,
and NH4Br (there are others). The insoluble bromides by Rule 3 would be AgBr, PbBr2, and
Hg2Br2. Similar reasoning is used for the other parts to this problem.
Sulfates: Na2SO4, K2SO4, and (NH4)2SO4 (and others) would be soluble, and BaSO4, CaSO4,
and PbSO4 (or Hg2SO4) would be insoluble.
Hydroxides: NaOH, KOH, Ca(OH)2 (and others) would be soluble, and Al(OH)3, Fe(OH)3,
and Cu(OH)2 (and others) would be insoluble.
Phosphates: Na3PO4, K3PO4, (NH4)3PO4 (and others) would be soluble, and Ag3PO4,
Ca3(PO4)2, and FePO4 (and others) would be insoluble.
Lead: PbCl2, PbBr2, PbI2, Pb(OH)2, PbSO4, and PbS (and others) would be insoluble.
Pb(NO3)2 would be a soluble Pb2+ salt.
18.
Pb(NO3)2(aq) + 2 KI(aq) → PbI2(s) + 2 KNO3(aq)
(formula equation)
Pb2+(aq) + 2 NO3−(aq) + 2 K+(aq) + 2 I−(aq) → PbI2(s) + 2 K+(aq) + 2 NO3−(aq)
(complete ionic equation)
The 1.0 mol of Pb2+ ions would react with the 2.0 mol of I− ions to form 1.0 mol of the PbI2
precipitate. Even though the Pb2+ and I− ions are removed, the spectator ions K+ and NO3− are
still present. The solution above the precipitate will conduct electricity because there are
plenty of charge carriers present in solution.
19.
The Brønsted-Lowry definitions are best for our purposes. An acid is a proton donor, and a
base is a proton acceptor. A proton is an H+ ion. Neutral hydrogen has 1 electron and 1
proton, so an H+ ion is just a proton. An acid-base reaction is the transfer of an H+ ion (a
proton) from an acid to a base.
20.
The acid is a diprotic acid (H2A), meaning that it has two H+ ions in the formula to donate to
a base. The reaction is H2A(aq) + 2 NaOH(aq) → 2 H2O(l) + Na2A(aq), where A2− is what is
left over from the acid formula when the two protons (H+ ions) are reacted.
CHAPTER 4
SOLUTION STOICHIOMETRY
95
For the HCl reaction, the base has the ability to accept two protons. The most common
examples are Ca(OH)2, Sr(OH)2, and Ba(OH)2. A possible reaction would be 2 HCl(aq) +
Ca(OH)2(aq) → 2 H2O(l) + CaCl2(aq).
21.
a. The species reduced is the element that gains electrons. The reducing agent causes reduction to occur by itself being oxidized. The reducing agent generally refers to the
entire formula of the compound/ion that contains the element oxidized.
b. The species oxidized is the element that loses electrons. The oxidizing agent causes
oxidation to occur by itself being reduced. The oxidizing agent generally refers to the
entire formula of the compound/ion that contains the element reduced.
c. For simple binary ionic compounds, the actual charges on the ions are the same as the
oxidation states. For covalent compounds, nonzero oxidation states are imaginary charges
the elements would have if they were held together by ionic bonds (assuming the bond is
between two different nonmetals). Nonzero oxidation states for elements in covalent
compounds are not actual charges. Oxidation states for covalent compounds are a
bookkeeping method to keep track of electrons in a reaction.
22.
Reference the Problem Solving Strategy box in Section 4.10 of the text for the steps involved
in balancing redox reactions by oxidation states. The key to the oxidation states method is to
balance the electrons gained by the species reduced with the number of electrons lost from
the species oxidized. This is done by assigning oxidation states and, from the change in
oxidation states, determining the coefficients necessary to balance electrons gained with
electrons lost. After the loss and gain of electrons is balanced, the remainder of the equation
is balanced by inspection.
Exercises
Aqueous Solutions: Strong and Weak Electrolytes
23.
a. NaBr(s) → Na+(aq) + Br-(aq)
b. MgCl2(s) → Mg2+(aq) + 2 Cl−(aq)
-
Na
Br
+
-
Na
+
Br
-
Na
Br
2+
Cl
-
Mg
-
Cl
Cl
+
-
Your drawing should show equal
number of Na+ and Br- ions.
2+
Mg
-
Cl
-
Cl
-
Cl
2+
Mg
Your drawing should show twice the
number of Cl− ions as Mg2+ ions.
96
CHAPTER 4
c. Al(NO3)3(s) → Al3+(aq) + 3 NO3−(aq)
Al3
+
NO 3
NO3
-
NO 3
-
NO 3
-
NO3
NO3
Al3
-
+
NO3
Al3
d. (NH4)2SO4(s) → 2 NH4+(aq) + SO42−(aq)
-
SO 4 2
+
NO3
NO3
SOLUTION STOICHIOMETRY
NH 4
-
+
-
SO 4 2
-
NH 4
NH 4
NH4
-
+
+
+
SO4 2
NH 4
NH 4
-
+
+
For e-i, your drawings should show equal numbers of the cations and anions present because
each salt is a 1 : 1 salt. The ions present are listed in the following dissolution reactions.
e. NaOH(s) → Na+(aq) + OH−(aq)
f.
g. KMnO4(s) → K+(aq) + MnO4− (aq)
h. HClO4(aq) → H+(aq) + ClO4−(aq)
i.
24.
FeSO4(s) → Fe2+(aq) + SO42−(aq)
NH4C2H3O2(s) → NH4+(aq) + C2H3O2−(aq)
a. Ba(NO3)2(aq) → Ba2+(aq) + 2 NO3−(aq);
present in Ba(NO3)2(aq).
picture iv represents the Ba2+ and NO3− ions
b. NaCl(aq) → Na+(aq) + Cl−(aq); picture ii represents NaCl(aq).
c. K2CO3(aq) → 2 K+(aq) + CO32−(aq); picture iii represents K2CO3(aq).
d. MgSO4(aq) → Mg2+(aq) + SO42−(aq); picture i represents MgSO4(aq).
HNO3(aq) → H+(aq) + NO3−(aq). Picture ii best represents the strong acid HNO3. Strong
acids are strong electrolytes. HC2H3O2 only partially dissociates in water; acetic acid is a
weak electrolyte. None of the pictures represent weak electrolyte solutions; they all are
representations of strong electrolytes.
25.
CaCl2(s) → Ca2+(aq) + 2 Cl−(aq)
26.
MgSO4(s) → Mg2+(aq) + SO42−(aq); NH4NO3(s) → NH4+(aq) + NO3−(aq)
Solution Concentration: Molarity
27.
a. 5.623 g NaHCO3 ×
M=
1 mol NaHCO 3
= 6.693 × 10 −2 mol NaHCO3
84.01 g NaHCO 3
6.693 × 10 −2 mol
1000 mL
= 0.2677 M NaHCO3
×
250.0 mL
L
CHAPTER 4
SOLUTION STOICHIOMETRY
b. 0.1846 g K2Cr2O7 ×
M=
500.0 × 10
c. 0.1025 g Cu ×
M=
1 mol K 2 Cr2 O 7
= 6.275 × 10 −4 mol K2Cr2O7
294.20 g K 2 Cr2 O 7
6.275 × 10 −4 mol
−3
97
L
= 1.255 × 10 −3 M K2Cr2O7
1 mol Cu
= 1.613 × 10 −3 mol Cu = 1.613 × 10 −3 mol Cu2+
63.55 g Cu
1.613 × 10 −3 mol Cu 2+
1000 mL
= 8.065 × 10 −3 M Cu2+
×
200.0 mL
L
28.
75.0 mL ×
29.
a.
1.3 mol
0.79 g
1 mol
×
= 1.3 mol C2H5OH; molarity =
= 5.2 M C2H5OH
0.250 L
mL
46.07 g
M Ca ( NO3 ) 2 =
0.100 mol Ca ( NO 3 ) 2
= 1.00 M
0.100 L
Ca(NO3)2(s) → Ca2+(aq) + 2 NO3−(aq); M Ca 2 + = 1.00 M; M NO − = 2(1.00) = 2.00 M
3
b.
M Na 2SO 4 =
2.5 mol Na 2 SO 4
1.25 L
= 2.0 M
Na2SO4(s) → 2 Na+(aq) + SO42−(aq); M Na + = 2(2.0) = 4.0 M ; M SO
c. 5.00 g NH4Cl ×
M NH 4Cl =
= 2.0 M
1 mol NH 4 Cl
= 0.0935 mol NH4Cl
53.49 g NH 4 Cl
0.0935 mol NH 4 Cl
= 0.187 M
0.5000 L
NH4Cl(s) → NH4+(aq) + Cl−(aq); M NH
d. 1.00 g K3PO4 ×
M K 3PO 4 =
2−
4
+
4
= M Cl − = 0.187 M
1 mol K 3 PO 4
= 4.71 × 10 −3 mol K3PO4
212.27 g
4.71 × 10 −3 mol
0.2500 L
= 0.0188 M
K3PO4(s) → 3 K+(aq) + PO43−(aq); M K + = 3(0.0188) = 0.0564 M; M PO3− = 0.0188 M
4
98
30.
CHAPTER 4
a.
M Na 3PO 4 =
SOLUTION STOICHIOMETRY
0.0200 mol
= 2.00 M
0.0100 L
Na3PO4(s) → 3 Na+(aq) + PO 34− (aq) ; M Na + = 3(2.00) = 6.00 M; M PO3− = 2.00 M
4
b.
M Ba ( NO3 ) 2 =
0.300 mol
= 0.500 M
0.6000 L
Ba(NO3)2(s) → Ba2+(aq) + 2 NO3−(aq); M Ba 2+ = 0.500 M; M NO − = 2(0.500) = 1.00 M
3
1 mol KCl
74.55 g KCl
= 0.0268 M
0.5000 L
1.00 g KCl ×
c.
M KCl =
KCl(s) → K+(aq) + Cl−(aq); M K + = M Cl − = 0.0268 M
1 mol ( NH 4 ) 2 SO 4
132.15 g
1.50 L
132 g ( NH 4 ) 2 SO 4 ×
d.
M ( NH 4 ) 2 SO 4 =
= 0.666 M
(NH4)2SO4(s) → 2 NH4+(aq) + SO42−(aq)
M NH
31.
+
4
= 2(0.666) = 1.33 M; M SO 2− = 0.666 M
4
 mol 
3+
−
Mol solute = volume (L) × molarity 
 ; AlCl3(s) → Al (aq) + 3 Cl (aq)
 L 
Mol Cl− = 0.1000 L ×
0.30 mol AlCl3
3 mol Cl −
= 9.0 × 10 −2 mol Cl−
×
L
mol AlCl3
MgCl2(s) → Mg2+(aq) + 2 Cl− (aq)
Mol Cl− = 0.0500 L ×
0.60 mol MgCl 2
2 mol Cl −
= 6.0 × 10 −2 mol Cl−
×
L
mol MgCl 2
NaCl(s) → Na+(aq) + Cl− (aq)
Mol Cl− = 0.2000 L ×
0.40 mol NaCl 1 mol Cl −
= 8.0 × 10 −2 mol Cl−
×
L
mol NaCl
100.0 mL of 0.30 M AlCl3 contains the most moles of Cl− ions.
32.
NaOH(s) → Na+(aq) + OH−(aq), 2 total mol of ions (1 mol Na+ and 1 mol Cl−) per
mol NaOH.
CHAPTER 4
SOLUTION STOICHIOMETRY
0.1000 L ×
99
0.100 mol NaOH
2 mol ions
×
= 2.0 × 10 −2 mol ions
L
mol NaOH
BaCl2(s) → Ba2+(aq) + 2 Cl−(aq), 3 total mol of ions per mol BaCl2.
0.0500 L ×
0.200 mol
3 mol ions
×
= 3.0 × 10 −2 mol ions
L
mol BaCl 2
Na3PO4(s) → 3 Na+(aq) + PO43−(aq), 4 total mol of ions per mol Na3PO4.
0.0750 L ×
0.150 mol Na 3 PO 4
4 mol ions
×
= 4.50 × mol 10 −2 ions
L
mol Na 3 PO 4
75.0 mL of 0.150 M Na3PO4 contains the largest number of ions.
33.
Molar mass of NaOH = 22.99 + 16.00 + 1.008 = 40.00 g/mol
Mass NaOH = 0.2500 L ×
34.
10. g AgNO3 ×
35.
a. 2.00 L ×
0.400 mol NaOH
40.00 g NaOH
×
= 4.00 g NaOH
L
mol NaOH
1 mol AgNO3
1L
×
= 0.24 L = 240 mL
169.9 g
0.25 mol AgNO3
0.250 mol NaOH
40.00 g NaOH
×
= 20.0 g NaOH
L
mol NaOH
Place 20.0 g NaOH in a 2-L volumetric flask; add water to dissolve the NaOH, and fill to
the mark with water, mixing several times along the way.
b. 2.00 L ×
0.250 mol NaOH
1 L stock
×
= 0.500 L
L
1.00 mol NaOH
Add 500. mL of 1.00 M NaOH stock solution to a 2-L volumetric flask; fill to the mark
with water, mixing several times along the way.
c. 2.00 L ×
0.100 mol K 2 CrO 4
194.20 g K 2 CrO 4
×
= 38.8 g K2CrO4
L
mol K 2 CrO 4
Similar to the solution made in part a, instead using 38.8 g K2CrO4.
d. 2.00 L ×
0.100 mol K 2 CrO 4
1 L stock
= 0.114 L
×
L
1.75 mol K 2 CrO 4
Similar to the solution made in part b, instead using 114 mL of the 1.75 M K2CrO4
stock solution.
100
36.
CHAPTER 4
0.50 mol H 2 SO 4
L
a. 1.00 L solution ×
SOLUTION STOICHIOMETRY
= 0.50 mol H2SO4
1L
= 2.8 × 10 −2 L conc. H2SO4 or 28 mL
18 mol H 2 SO 4
0.50 mol H2SO4 ×
Dilute 28 mL of concentrated H2SO4 to a total volume of 1.00 L with water. The resulting
1.00 L of solution will be a 0.50 M H2SO4 solution.
b. We will need 0.50 mol HCl.
0.50 mol HCl ×
1L
= 4.2 × 10 −2 L = 42 mL
12 mol HCl
Dilute 42 mL of concentrated HCl to a final volume of 1.00 L.
c. We need 0.50 mol NiCl2.
0.50 mol NiCl2 ×
1 mol NiCl 2 • 6H 2 O
237.69 g NiCl 2 • 6H 2 O
×
mol NiCl 2
mol NiCl 2 • 6H 2 O
= 118.8 g NiCl2•6H2O ≈ 120 g
Dissolve 120 g NiCl2•6H2O in water, and add water until the total volume of the solution
is 1.00 L.
d. 1.00 L ×
0.50 mol HNO3
= 0.50 mol HNO3
L
0.50 mol HNO3 ×
1L
= 0.031 L = 31 mL
16 mol HNO3
Dissolve 31 mL of concentrated reagent in water. Dilute to a total volume of 1.00 L.
e. We need 0.50 mol Na2CO3.
0.50 mol Na2CO3 ×
105.99 g Na 2 CO 3
mol
= 53 g Na2CO3
Dissolve 53 g Na2CO3 in water, dilute to 1.00 L.
37.
10.8 g (NH4)2SO4 ×
Molarity =
1 mol
= 8.17 × 10 −2 mol (NH4)2SO4
132.15 g
8.17 × 10 −2 mol
1000 mL
= 0.817 M (NH4)2SO4
×
100.0 mL
L
CHAPTER 4
SOLUTION STOICHIOMETRY
101
Moles of (NH4)2SO4 in final solution:
10.00 × 10-3 L ×
0.817 mol
= 8.17 × 10 −3 mol
L
8.17 × 10 −3 mol
1000 mL
= 0.136 M (NH4)2SO4
×
(10.00 + 50.00) mL
L
Molarity of final solution =
(NH4)2SO4(s) → 2 NH4+(aq) + SO42−(aq); M NH
38.
Molarity =
+
4
= 2(0.136) = 0.272 M; M SO 2− = 0.136 M
4
total mol HNO3
; total volume = 0.05000 L + 0.10000 L = 0.15000 L
total volume
Total mol HNO3 = 0.05000 L ×
0.100 mol HNO3
0.200 mol HNO3
+ 0.10000 L ×
L
L
Total mol HNO3 = 5.00 × 10−3 mol + 2.00 × 10−2 mol = 2.50 × 10−2 mol HNO3
Molarity =
2.50 × 10 −2 mol HNO3
= 0.167 M HNO3
0.15000 L
As expected, the molarity of HNO3 is between 0.100 M and 0.200 M.
39.
3.0 mol Na 2 CO 3
L
Mol Na2CO3 = 0.0700 L ×
= 0.21 mol Na2CO3
Na2CO3(s) → 2 Na+(aq) + CO32−(aq); mol Na+ = 2(0.21 mol) = 0.42 mol
Mol NaHCO3 = 0.0300 L ×
1.0 mol NaHCO 3
L
= 0.030 mol NaHCO3
NaHCO3(s) → Na+(aq) + HCO3−(aq); mol Na+ = 0.030 mol
M Na + =
40.
total mol Na +
total volume
=
0.42 mol + 0.030 mol
0.45 mol
=
= 4.5 M Na+
0.0700 L + 0.0300 L
0.1000 L
Mol CoCl2 = 0.0500 L ×
0.250 mol CoCl 2
= 0.0125 mol
L
Mol NiCl2 = 0.0250 L ×
0.350 mol NiCl 2
= 0.00875 mol
L
Both CoCl2 and NiCl2 are soluble chloride salts by the solubility rules. A 0.0125-mol
aqueous sample of CoCl2 is actually 0.0125 mol Co2+ and 2(0.0125 mol) = 0.0250 mol Cl−.
A 0.00875-mol aqueous sample of NiCl2 is actually 0.00875 mol Ni2+ and 2(0.00875) =
0.0175 mol Cl−. The total volume of solution that these ions are in is 0.0500 L + 0.0250 L =
0.0750 L.
102
CHAPTER 4
M Co 2 + =
M Cl − =
41.
SOLUTION STOICHIOMETRY
0.0125 mol Co 2+
0.00875 mol Ni 2+
= 0.167 M ; M Ni 2 + =
= 0.117 M
0.0750 L
0.0750 L
0.0250 mol Cl − + 0.0175 mol Cl −
= 0.567 M
0.0750 L
Stock solution =
10.0 mg
10.0 × 10 −3 g
2.00 × 10 −5 g steroid
=
=
500.0 mL
500.0 mL
mL
100.0 × 10 −6 L stock ×
1000 mL 2.00 × 10 −5 g steroid
= 2.00 × 10 −6 g steroid
×
L
mL
This is diluted to a final volume of 100.0 mL.
2.00 × 10 −6 g steroid
1000 mL
1 mol steroid
= 5.94 × 10 −8 M steroid
×
×
100.0 mL
L
336.43 g steroid
42.
Stock solution:
1.584 g Mn2+ ×
1 mol Mn 2+
54.94 g Mn
2+
= 2.883 × 10 −2 mol Mn2+
2.833 × 10 −2 mol Mn 2+
= 2.883 × 10 −2 M
1.000 L
Molarity =
Solution A:
50.00 mL ×
1L
2.833 × 10 −2 mol
×
1000 mL
L
Molarity =
1.442 × 10 −3 mol
1000.0 mL
×
= 1.442 × 10 −3 mol Mn2+
1000 mL
= 1.442 × 10 −3 M
1L
Solution B:
10.0 mL ×
Molarity =
1L
1.442 × 10 −3 mol
×
1000 mL
L
1.442 × 10 −5 mol
0.2500 L
= 1.442 × 10 −5 mol Mn2+
= 5.768 × 10 −5 M
Solution C:
10.00 × 10-3 L ×
Molarity =
5.768 × 10 −5 mol
L
5.768 × 10 −7 mol
0.5000 L
= 5.768 × 10 −7 mol Mn2+
= 1.154 × 10 −6 M
CHAPTER 4
SOLUTION STOICHIOMETRY
103
Precipitation Reactions
43.
The solubility rules referenced in the following answers are outlined in Table 4.1 of the text.
a. Soluble: Most nitrate salts are soluble (Rule 1).
b. Soluble: Most chloride salts are soluble except for Ag+, Pb2+, and Hg22+ (Rule 3).
c. Soluble: Most sulfate salts are soluble except for BaSO4, PbSO4, Hg2SO4, and CaSO4
(Rule 4.)
d. Insoluble: Most hydroxide salts are only slightly soluble (Rule 5).
Note: We will interpret the phrase “slightly soluble” as meaning insoluble and the phrase
“marginally soluble” as meaning soluble. So the marginally soluble hydroxides Ba(OH)2,
Sr(OH)2, and Ca(OH)2 will be assumed soluble unless noted otherwise.
e. Insoluble: Most sulfide salts are only slightly soluble (Rule 6). Again, “slightly soluble”
is interpreted as “insoluble” in problems like these.
f.
Insoluble: Rule 5 (see answer d).
g. Insoluble: Most phosphate salts are only slightly soluble (Rule 6).
44.
45.
The solubility rules referenced in the following answers are from Table 4.1 of the text. The
phrase “slightly soluble” is interpreted to mean insoluble, and the phrase “marginally soluble”
is interpreted to mean soluble.
a. Soluble (Rule 3)
b. Soluble (Rule 1)
c. Inoluble (Rule 4)
d. Soluble (Rules 2 and 3)
e. Insoluble (Rule 6)
f.
g. Insoluble (Rule 6)
h. Soluble (Rule 2)
Insoluble (Rule 5)
In these reactions, soluble ionic compounds are mixed together. To predict the precipitate,
switch the anions and cations in the two reactant compounds to predict possible products;
then use the solubility rules in Table 4.1 to predict if any of these possible products are
insoluble (are the precipitate). Note that the phrase “slightly soluble” in Table 4.1 is
interpreted to mean insoluble, and the phrase “marginally soluble” is interpreted to mean
soluble.
a. Possible products = FeCl2 and K2SO4; both salts are soluble, so no precipitate forms.
b. Possible products = Al(OH)3 and Ba(NO3)2; precipitate = Al(OH)3(s)
c. Possible products = CaSO4 and NaCl; precipitate = CaSO4(s)
d. Possible products = KNO3 and NiS; precipitate = NiS(s)
46.
Use Table 4.1 to predict the solubility of the possible products.
a. Possible products = Hg2SO4 and Cu(NO3)2; precipitate = Hg2SO4
b. Possible products = NiCl2 and Ca(NO3)2; both salts are soluble so no precipitate forms.
104
CHAPTER 4
SOLUTION STOICHIOMETRY
c. Possible products = KI and MgCO3; precipitate = MgCO3
d. Possible products = NaBr and Al2(CrO4)3; precipitate = Al2(CrO4)3
47.
For the following answers, the balanced formula equation is first, followed by the complete
ionic equation, then the net ionic equation.
a. No reaction occurs since all possible products are soluble salts.
b. 2 Al(NO3)3(aq) + 3 Ba(OH)2(aq) → 2 Al(OH)3(s) + 3 Ba(NO3)2(aq)
2 Al3+(aq) + 6 NO3−(aq) + 3 Ba2+(aq) + 6 OH−(aq) →
2 Al(OH)3(s) + 3 Ba2+(aq) + 6 NO3−(aq)
3+
−
Al (aq) + 3 OH (aq) → Al(OH)3(s)
c. CaCl2(aq) + Na2SO4(aq) → CaSO4(s) + 2 NaCl(aq)
Ca2+(aq) + 2 Cl−(aq) + 2 Na+(aq) + SO42−(aq) → CaSO4(s) + 2 Na+(aq) + 2 Cl−(aq)
Ca2+(aq) + SO42−(aq) → CaSO4(s)
d. K2S(aq) + Ni(NO3)2(aq) → 2 KNO3(aq) + NiS(s)
2 K+(aq) + S2−(aq) + Ni2+(aq) + 2 NO3−(aq) → 2 K+(aq) + 2 NO3−(aq) + NiS(s)
Ni2+(aq) + S2−(aq) → NiS(s)
48.
a. Hg2(NO3)2(aq) + CuSO4(aq) → Hg2SO4(s) + Cu(NO3)2(aq)
Hg22+(aq) + 2 NO3−(aq) + Cu2+(aq) + SO42−(aq) → Hg2SO4(s) + Cu2+(aq) + 2 NO3−(aq)
Hg22+(aq) + SO42−(aq) → Hg2SO4(s)
b. No reaction occurs since both possible products are soluble.
c. K2CO3(aq) + MgI2(aq) → 2 KI(aq) + MgCO3(s)
2 K+(aq) + CO32−(aq) + Mg2+(aq) + 2I−(aq) → 2 K+(aq) + 2 I−(aq) + MgCO3(s)
Mg2+(aq) + CO32−(aq) → MgCO3(s)
d. 3 Na2CrO4(aq) + 2 Al(Br)3(aq) → 6 NaBr(aq) + Al2(CrO4)3(s)
6 Na+(aq) + 3 CrO42−(aq) + 2 Al3+(aq) + 6 Br−(aq) → 6 Na+(aq) + 6 Br−(aq) +
2 Al3+(aq) + 3 CrO42−(aq) → Al2(CrO4)3(s)
49.
Al2(CrO4)3(s)
a. When CuSO4(aq) is added to Na2S(aq), the precipitate that forms is CuS(s). Therefore,
Na+ (the gray spheres) and SO42− (the bluish green spheres) are the spectator ions.
CuSO4(aq) + Na2S(aq) → CuS(s) + Na2SO4(aq); Cu2+(aq) + S2−(aq) → CuS(s)
CHAPTER 4
SOLUTION STOICHIOMETRY
105
b. When CoCl2(aq) is added to NaOH(aq), the precipitate that forms is Co(OH)2(s).
Therefore, Na+ (the gray spheres) and Cl- (the green spheres) are the spectator ions.
CoCl2(aq) + 2 NaOH(aq) → Co(OH)2(s) + 2 NaCl(aq)
Co2+(aq) + 2 OH−(aq) → Co(OH)2(s)
c. When AgNO3(aq) is added to KI(aq), the precipitate that forms is AgI(s). Therefore, K+
(the red spheres) and NO3− (the blue spheres) are the spectator ions.
AgNO3(aq) + KI(aq) → AgI(s) + KNO3(aq); Ag+(aq) + I−(aq) → AgI(s)
50.
There are many acceptable choices for spectator ions. We will generally choose Na+ and
NO3− as the spectator ions because sodium salts and nitrate salts are usually soluble in water.
a. Fe(NO3)3(aq) + 3 NaOH(aq) → Fe(OH)3(s) + 3 NaNO3(aq)
b. Hg2(NO3)2(aq) + 2 NaCl(aq) → Hg2Cl2(s) + 2 NaNO3(aq)
c. Pb(NO3)2(aq) + Na2SO4(aq) → PbSO4(s) + 2 NaNO3(aq)
d. BaCl2(aq) + Na2CrO4(aq) → BaCrO4(s) + 2 NaCl(aq)
51.
a. (NH4)2SO4(aq) + Ba(NO3)2(aq) → 2 NH4NO3(aq) + BaSO4(s)
Ba2+(aq) + SO42−(aq) → BaSO4(s)
b. Pb(NO3)2(aq) + 2 NaCl(aq) → PbCl2(s) + 2 NaNO3(aq)
Pb2+(aq) + 2 Cl−(aq) → PbCl2(s)
c. Potassium phosphate and sodium nitrate are both soluble in water. No reaction occurs.
d. No reaction occurs because all possible products are soluble.
e. CuCl2(aq) + 2 NaOH(aq) → Cu(OH)2(s) + 2 NaCl(aq)
Cu2+(aq) + 2 OH−(aq) → Cu(OH)2(s)
52.
a. CrCl3(aq) + 3 NaOH(aq) → Cr(OH)3(s) + 3 NaCl(aq)
Cr3+(aq) + 3 OH−(aq) → Cr(OH)3(s)
b. 2 AgNO3(aq) + (NH4)2CO3(aq) → Ag2CO3(s) + 2 NH4NO3(aq)
2 Ag+(aq) + CO32−(aq) → Ag2CO3(s)
c. CuSO4(aq) + Hg2(NO3)2(aq) → Cu(NO3)2(aq) + Hg2SO4(s)
106
CHAPTER 4
SOLUTION STOICHIOMETRY
Hg22+(aq) + SO42−(aq) → Hg2SO4(s)
d. No reaction occurs because all possible products (SrI2 and KNO3) are soluble.
53.
Because a precipitate formed with Na2SO4, the possible cations are Ba2+, Pb2+, Hg22+, and
Ca2+ (from the solubility rules). Because no precipitate formed with KCl, Pb2+ and Hg22+
cannot be present. Because both Ba2+ and Ca2+ form soluble chlorides and soluble
hydroxides, both these cations could be present. Therefore, the cations could be Ba2+ and Ca2+
(by the solubility rules in Table 4.1). For students who do a more rigorous study of solubility,
Sr2+ could also be a possible cation (it forms an insoluble sulfate salt, whereas the chloride
and hydroxide salts of strontium are soluble).
54.
Because no precipitates formed upon addition of NaCl or Na2SO4, we can conclude that Hg22+
and Ba2+ are not present in the sample because Hg2Cl2 and BaSO4 are insoluble salts.
However, Mn2+ may be present since Mn2+ does not form a precipitate with either NaCl or
Na2SO4. A precipitate formed with NaOH; the solution must contain Mn2+ because it forms
a precipitate with OH− [Mn(OH)2(s)].
55.
2 AgNO3(aq) + Na2CrO4(aq) → Ag2CrO4(s) + 2 NaNO3(aq)
0.0750 L ×
56.
0.100 mol AgNO3 1 mol Na 2 CrO 4
161.98 g Na 2 CrO 4
×
×
= 0.607 g Na2CrO4
L
2 mol AgNO3
mol Na 2 CrO 4
2 Na3PO4(aq) + 3 Pb(NO3)2(aq) → Pb3(PO4)2(s) + 6 NaNO3(aq)
0.1500 L ×
57.
0.250 mol Pb( NO 3 ) 2
2 mol Na 3PO 4
1 L Na 3 PO 4
×
×
= 0.250 L
L
3 mol Pb( NO 3 ) 2
0.100 mol Na 3 PO 4
A1(NO3)3(aq) + 3 KOH(aq) → Al(OH)3(s) + 3 KNO3(aq)
= 250. mL Na3PO4
Assuming Al(NO3)3 is limiting:
0.200 mol Al( NO 3 ) 3
1 mol Al(OH) 3
78.00 g Al(OH) 3
×
×
mol Al( NO 3 ) 3
mol Al(OH) 3
L
= 0.780 g Al(OH)3
Assuming KOH is limiting:
0.0500 L ×
0.2000 L ×
1 mol Al(OH) 3
78.00 g Al(OH) 3
0.100 mol KOH
×
×
= 0.520 g Al(OH)3
3 mol KOH
mol Al(OH) 3
L
Because KOH produces the smaller mass of the Al(OH)3 precipitate, KOH is the limiting
reagent and 0.520 g Al(OH)3 can form.
58.
The balanced equation is 3 BaCl2(aq) + Fe2(SO4)3(aq) → 3 BaSO4(s) + 2 FeCl3(aq).
100.0 mL BaCl2 ×
0.100 mol BaCl 2
3 mol BaSO 4
233.4 g BaSO 4
1L
×
×
×
1000 mL
L
3 mol BaCl2
mol BaSO 4
= 2.33 g BaSO4
CHAPTER 4
SOLUTION STOICHIOMETRY
100.0 mL Fe2(SO4)3 ×
107
0.100 mol Fe 2 (SO 4 ) 3
3 mol BaSO 4
1L
×
×
1000 mL
L
mol Fe 2 (SO 4 ) 3
233.4 g BaSO 4
×
= 7.00 g BaSO4
mol BaSO 4
The BaCl2 reagent produces the smaller quantity of the BaSO4 precipitate, so BaCl2 is
limiting and 2.33 g BaSO4 can form.
59.
The reaction is AgNO3(aq) + NaBr(aq) → AgBr(s) + NaNO3(aq).
Assuming AgNO3 is limiting:
100.0 mL AgNO3 ×
0.150 mol AgNO3
1L
1 mol AgBr
187.8 g AgBr
×
×
×
1000 mL
L AgNO3
mol AgNO3
mol AgBr
= 2.82 g AgBr
Assuming NaBr is limiting:
20.0 mL NaBr ×
1L
1.00 mol NaBr 1 mol AgBr 187.8 g AgBr
×
×
×
mol NaBr
1000 mL
L NaBr
mol AgBr
= 3.76 g AgBr
The AgNO3 reagent produces the smaller quantity of AgBr, so AgNO3 is limiting and 2.82 g
AgBr can form.
60.
2 AgNO3(aq) + CaCl2(aq) → 2 AgCl(s) + Ca(NO3)2(aq)
0.1000 L ×
0.20 mol AgNO3
2 mol AgCl
143.4 g AgCl
×
×
= 2.9 g AgCl
2 mol AgNO3
mol AgCl
L
0.1000 L ×
0.15 mol CaCl 2 2 mol AgCl 143.4 g AgCl
×
×
= 4.3 g AgCl
mol CaCl 2
mol AgCl
L
AgNO3 is limiting (it produces the smaller mass of AgCl) and 2.9 g AgCl can form.
The net ionic equation is Ag+(aq) + Cl−(aq) → AgCl(s). The ions remaining in solution are
the unreacted Cl− ions and the spectator ions NO3− and Ca2+ (all Ag+ is used up in forming
AgCl). The moles of each ion present initially (before reaction) can be easily determined
from the moles of each reactant. We have 0.1000 L(0.20 mol AgNO3/L) = 0.020 mol
AgNO3, which dissolves to form 0.020 mol Ag+ and 0.020 mol NO3−. We also have
0.1000 L(0.15 mol CaCl2/L) = 0.015 mol CaCl2, which dissolves to form 0.015 mol Ca2+ and
2(0.015) = 0.030 mol Cl−. To form the 2.9 g of AgCl precipitate, 0.020 mol Ag+ will react
with 0.020 mol of Cl− to form 0.020 mol AgCl (which has a mass of 2.9 g).
Mol unreacted Cl− = 0.030 mol Cl− initially − 0.020 mol Cl− reacted
Mol unreacted Cl− = 0.010 mol Cl−
108
CHAPTER 4
M Cl− =
SOLUTION STOICHIOMETRY
0.010 mol Cl −
0.010 mol Cl −
= 0.050 M Cl−
=
total volume
0.1000 L + 0.1000 L
The molarities of the spectator ions are:
0.020 mol NO 3
0.2000 L
61.
−
= 0.10 M NO3−;
0.015 mol Ca 2+
= 0.075 M Ca2+
0.2000 L
a. The balanced reaction is 2 KOH(aq) + Mg(NO3)2(aq) → Mg(OH)2(s) + 2 KNO3(aq).
b. The precipitate is magnesium hydroxide.
c. Assuming KOH is limiting:
0.1000 L KOH ×
0.200 mol KOH 1 mol Mg (OH) 2 58.33 g Mg (OH) 2
×
×
L KOH
2 mol KOH
mol Mg (OH) 2
= 0.583 g Mg(OH)2
Assuming Mg(NO3)2 is limiting:
0.1000 L Mg(NO3)2 ×
0.200 mol Mg(NO 3 ) 2
1 mol Mg(OH) 2
×
mol Mg(NO 3 ) 2
L Mg(NO 3 ) 2
×
58.33 g Mg(OH) 2
= 1.17 g Mg(OH)2
mol Mg(OH) 2
The KOH reagent is limiting because it produces the smaller quantity of the Mg(OH)2
precipitate. So 0.583 g Mg(OH)2 can form.
d. The net ionic equation for this reaction is Mg2+(aq) + 2 OH−(aq) → Mg(OH)2(s).
Because KOH is the limiting reagent, all of the OH− is used up in the reaction. So M OH −
= 0 M. Note that K+ is a spectator ion, so it is still present in solution after precipitation
was complete. Also present will be the excess Mg2+ and NO3− (the other spectator ion).
Total Mg2+ = 0.1000 L Mg(NO3)2 ×
0.200 mol Mg ( NO 3 ) 2
1 mol Mg 2+
×
L Mg ( NO 3 ) 2
mol Mg ( NO 3 ) 2
= 0.0200 mol Mg2+
Mol Mg2+ reacted = 0.1000 L KOH ×
0.200 mol KOH 1 mol Mg(NO 3 ) 2
×
L KOH
2 mol KOH
×
M Mg 2 + =
1 mol Mg 2+
= 0.0100 mol Mg2+
mol Mg ( NO 3 ) 2
(0.0200 − 0.0100) mol Mg 2+
mol excess Mg 2+
=
= 5.00 × 10−2 M Mg2+
0.1000 L + 0.1000 L
total volume
CHAPTER 4
SOLUTION STOICHIOMETRY
109
The spectator ions are K+ and NO3−. The moles of each are:
mol K+ = 0.1000 L KOH ×
0.200 mol KOH
1 mol K +
= 0.0200 mol K+
×
L KOH
mol KOH
mol NO3− = 0.1000 L Mg(NO3)2 ×
0.200 mol Mg ( NO 3 ) 2
2 mol NO 3 −
×
L Mg ( NO 3 ) 2
mol Mg ( NO 3 ) 2
= 0.0400 mol NO3−
The concentrations are:
0.0400 mol NO 3 −
0.0200 mol K +
= 0.100 M K+;
= 0.200 M NO3−
0.2000 L
0.2000 L
62.
a. Cu(NO3)2(aq) + 2 KOH(aq) → Cu(OH)2(s) + 2 KNO3(aq)
Solution A contains 2.00 L × 2.00 mol/L = 4.00 mol Cu(NO3)2, and solution B contains
2.00 L × 3.00 mol/L = 6.00 mol KOH. In the picture in the problem, we have 4 formula
units of Cu(NO3)2 (4 Cu2+ ions and 8 NO3− ions) and 6 formula units of KOH (6 K+ ions
and 6 OH− ions). With 4 Cu2+ ions and 6 OH− ions present, OH− is limiting (when all 6
molecules of OH− react, we only need 3 of the 4 Cu2+ ions to react with all of the OH−
present). After reaction, one Cu2+ ion remains as 3 Cu(OH)2(s) formula units form as
precipitate. The following drawing summarizes the ions that remain in solution and the
relative amount of precipitate that forms. Note that K+ and NO3− ions are spectator ions.
In the drawing, V1 is the volume of solution A or B, and V2 is the volume of the
combined solutions, with V2 = 2V1. The drawing exaggerates the amount of precipitate
that would actually form.
V2
NO3
V1
K
-
+
K
NO3
+
NO 3
+
NO 3
-
2+
NO3
-
K
-
K
-
Cu
NO3
+
NO 3
-
K
-
K
+
+
NO 3
-
Cu(OH)2
Cu(OH)2
Cu(OH)2
b. The spectator ion concentrations will be one-half the original spectator ion concentrations
in the individual beakers because the volume was doubled. Or using moles, M K + =
8.00 mol NO 3
6.00 mol K +
= 1.50 M and M NO − =
3
4.00 L
4.00 L
−
= 2.00 M. The concentration of
OH− ions will be zero because OH− is the limiting reagent. From the drawing, the number
of Cu2+ ions will decrease by a factor of four as the precipitate forms. Because the
110
CHAPTER 4
SOLUTION STOICHIOMETRY
volume of solution doubled, the concentration of Cu2+ ions will decrease by a factor of
eight after the two beakers are mixed:
1
M Cu + = 2.00 M   = 0.250 M
8
Alternately, one could certainly use moles to solve for M Cu 2 + :
Mol Cu2+ reacted = 2.00 L ×
3.00 mol OH − 1 mol Cu 2+
= 3.00 mol Cu2+ reacted
×
L
2 mol OH −
Mol Cu2+ present initially = 2.00 L ×
2.00 mol Cu 2+
= 4.00 mol Cu2+ present initially
L
Excess Cu2+ present after reaction = 4.00 mol − 3.00 mol = 1.00 mol Cu2+ excess
M Cu 2+ =
1.00 mol Cu 2+
= 0.250 M
2.00 L + 2.00 L
Mass of precipitate = 6.00 mol KOH ×
63.
1 mol Cu (OH) 2 97.57 g Cu (OH) 2
×
2 mol KOH
mol Cu (OH) 2
= 293 g Cu(OH)2
M2SO4(aq) + CaCl2(aq) → CaSO4(s) + 2 MCl(aq)
1.36 g CaSO4 ×
1 mol CaSO 4
1 mol M 2SO 4
×
= 9.99 × 10 −3 mol M2SO4
136.15 g CaSO 4
mol CaSO 4
From the problem, 1.42 g M2SO4 was reacted, so:
molar mass =
1.42 g M 2SO 4
9.99 × 10 −3 mol M 2SO 4
= 142 g/mol
142 u = 2(atomic mass M) + 32.07 + 4(16.00), atomic mass M = 23 u
From periodic table, M = Na (sodium).
64.
a. Na+, NO3−, Cl−, and Ag+ ions are present before any reaction occurs. The excess Ag+
added will remove all of the Cl− ions present. Therefore, Na+, NO3−, and the excess Ag+
ions will all be present after precipitation of AgCl is complete.
b. Ag+(aq) + Cl−(aq) → AgCl(s)
c. Mass NaCl = 0.641 g AgCl ×
Mass % NaCl =
1 mol AgCl 1 mol Cl −
1 mol NaCl
58.44 g
×
×
×
−
143.4 g
mol AgCl
mol NaCl
mol Cl
= 0.261 g NaCl
0.261 g NaCl
× 100 = 17.4% NaCl
1.50 g mixture
CHAPTER 4
SOLUTION STOICHIOMETRY
111
Acid-Base Reactions
65.
All the bases in this problem are ionic compounds containing OH-. The acids are either
strong or weak electrolytes. The best way to determine if an acid is a strong or weak
electrolyte is to memorize all the strong electrolytes (strong acids). Any other acid you
encounter that is not a strong acid will be a weak electrolyte (a weak acid), and the formula
should be left unaltered in the complete ionic and net ionic equations. The strong acids to
recognize are HCl, HBr, HI, HNO3, HClO4, and H2SO4. For the following answers, the order
of the equations are formula, complete ionic, and net ionic.
a. 2 HClO4(aq) + Mg(OH)2(s) → 2 H2O(l) + Mg(ClO4)2(aq)
2 H+(aq) + 2 ClO4−(aq) + Mg(OH)2(s) → 2 H2O(l) + Mg2+(aq) + 2 ClO4−(aq)
2 H+(aq) + Mg(OH)2(s) → 2 H2O(l) + Mg2+(aq)
b. HCN(aq) + NaOH(aq) → H2O(l) + NaCN(aq)
HCN(aq) + Na+(aq) + OH−(aq) → H2O(l) + Na+(aq) + CN−(aq)
HCN(aq) + OH−(aq) → H2O(l) + CN−(aq)
c. HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
H+(aq) + Cl−(aq) + Na+(aq) + OH−(aq) → H2O(l) + Na+(aq) + Cl−(aq)
H+(aq) + OH−(aq) → H2O(l)
66.
a. 3 HNO3(aq) + Al(OH)3(s) → 3 H2O(l) + Al(NO3)3(aq)
3 H+(aq) + 3 NO3−(aq) + Al(OH)3(s) → 3 H2O(l) + Al3+(aq) + 3 NO3−(aq)
3 H+(aq) + Al(OH)3(s) → 3 H2O(l) + Al3+(aq)
b. HC2H3O2(aq) + KOH(aq) → H2O(l) + KC2H3O2(aq)
HC2H3O2(aq) + K+(aq) + OH−(aq) → H2O(l) + K+(aq) + C2H3O2−(aq)
HC2H3O2(aq) + OH−(aq) → H2O(l) + C2H3O2−(aq)
c. Ca(OH)2(aq) + 2 HCl(aq) → 2 H2O(l) + CaCl2(aq)
Ca2+(aq) + 2 OH−(aq) + 2 H+(aq) + 2 Cl−(aq) → 2 H2O(l) + Ca2+(aq) + 2 Cl−(aq)
2 H+(aq) + 2 OH−(aq) → 2 H2O(l) or H+(aq) + OH−(aq) → H2O(l)
67.
All the acids in this problem are strong electrolytes (strong acids). The acids to recognize as
strong electrolytes are HCl, HBr, HI, HNO3, HClO4, and H2SO4.
112
CHAPTER 4
SOLUTION STOICHIOMETRY
a. KOH(aq) + HNO3(aq) → H2O(l) + KNO3(aq)
b. Ba(OH)2(aq) + 2 HCl(aq) → 2 H2O(l) + BaCl2(aq)
c. 3 HClO4(aq) + Fe(OH)3(s) → 3 H2O(l) + Fe(ClO4)3(aq)
d. AgOH(s) + HBr(aq) → AgBr(s) + H2O(l)
e. Sr(OH)2(aq) + 2 HI(aq) → 2 H2O(l) + SrI2(aq)
68.
a. Perchloric acid plus potassium hydroxide is a possibility.
HClO4(aq) + KOH(aq) → H2O(l) + KClO4(aq)
b. Nitric acid plus cesium hydroxide is a possibility.
HNO3(aq) + CsOH(aq) → H2O(l) + CsNO3(aq)
c. Hydroiodic acid plus calcium hydroxide is a possibility.
2 HI(aq) + Ca(OH)2(aq) → 2 H2O(l) + CaI2(aq)
69.
If we begin with 50.00 mL of 0.200 M NaOH, then:
50.00 × 10 −3 L ×
0.200 mol
= 1.00 × 10 −2 mol NaOH is to be neutralized
L
a. NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
1.00 × 10 −2 mol NaOH ×
1 mol HCl
1L
×
= 0.100 L or 100. mL
mol NaOH
0.100 mol
b. HNO3(aq) + NaOH(aq) → H2O(l) + NaNO3(aq)
1.00 × 10 −2 mol NaOH ×
1 mol HNO3
1L
×
= 6.67 × 10 −2 L or 66.7 mL
mol NaOH
0.150 mol HNO3
c. HC2H3O2(aq) + NaOH(aq) → H2O(l) + NaC2H3O2(aq)
1.00 × 10 −2 mol NaOH ×
70.
1 mol HC 2 H 3O 2
1L
×
= 5.00 × 10 −2 L
mol NaOH
0.200 mol HC 2 H 3O 2
= 50.0 mL
We begin with 25.00 mL of 0.200 M HCl or 25.00 × 10 −3 L × 0.200 mol/L
= 5.00 × 10 −3 mol HCl.
a. HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
5.00 × 10 −3 mol HCl ×
1 mol HCl
1L
×
= 5.00 × 10 −2 L or 50.0 mL
mol NaOH
0.100 mol NaOH
b. 2 HCl(aq) + Sr(OH)2(aq) → 2 H2O(l) + SrCl2(aq)
CHAPTER 4
SOLUTION STOICHIOMETRY
5.00 × 10 −3 mol HCl ×
113
1 mol Sr(OH) 2
1L
×
= 5.00 × 10 −2 L
2 mol HCl
0.0500 mol Sr(OH) 2
= 50.0 mL
c. HCl(aq) + KOH(aq) → H2O(l) + KCl(aq)
5.00 × 10 −3 mol HCl ×
71.
1 mol KOH
1L
×
= 2.00 × 10 −2 L = 20.0 mL
mol HCl
0.250 mol KOH
Ba(OH)2(aq) + 2 HCl(aq) → BaCl2(aq) + 2 H2O(l); H+(aq) + OH−(aq) → H2O(l)
75.0 × 10 −3 L ×
0.250 mol HCl
L
= 1.88 × 10 −2 mol HCl = 1.88 × 10 −2 mol H+
+ 1.88 × 10 −2 mol Cl−
225.0 × 10 −3 L ×
0.0550 mol Ba (OH) 2
= 1.24 × 10 −2 mol Ba(OH)2
L
= 1.24 × 10 −2 mol Ba2+ + 2.48 × 10 −2 mol OH−
The net ionic equation requires a 1 : 1 mole ratio between OH− and H+. The actual mole OH−
to mole H+ ratio is greater than 1 : 1, so OH− is in excess. Because 1.88 × 10 −2 mol OH− will
be neutralized by the H+, we have (2.48 − 1.88) × 10 −2 = 0.60 × 10 −2 mol OH− in excess.
M OH − =
72.
mol OH − excess
total volume
=
6.0 × 10 −3 mol OH −
= 2.0 × 10 −2 M OH−
0.0750 L + 0.2250 L
HCl and HNO3 are strong acids; Ca(OH)2 and RbOH are strong bases. The net ionic equation
that occurs is H+(aq) + OH−(aq) → H2O(l).
Mol H+ = 0.0500 L ×
0.100 mol HCl 1 mol H +
×
L
mol HCl
0.200 mol HNO 3
1 mol H +
= 0.00500 + 0.0200 = 0.0250 mol H+
×
L
mol HNO 3
+ 0.1000 L ×
Mol OH− = 0.5000 L ×
+ 0.2000 L ×
0.0100 mol Ca (OH) 2
2 mol OH −
×
L
mol Ca (OH) 2
0.100 mol RbOH 1 mol OH −
= 0.0100 + 0.0200 = 0.0300 mol OH−
×
L
mol RbOH
We have an excess of OH−, so the solution is basic (not neutral). The moles of excess OH− =
0.0300 mol OH− initially − 0.0250 mol OH− reacted (with H+) = 0.0050 mol OH− excess.
M OH − =
0.0050 mol OH −
0.0050 mol
= 5.9 × 10 −3 M
=
(0.0500 + 0.1000 + 0.5000 + 0.2000) L
0.8500 L
114
73.
CHAPTER 4
HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
0.106 mol NaOH
1 mol HCl
×
= 2.56 × 10 −3 mol HCl
L NaOH
mol NaOH
24.16 × 10 −3 L NaOH ×
Molarity of HCl =
74.
SOLUTION STOICHIOMETRY
2.56 × 10 −3 mol
25.00 × 10 −3 L
= 0.102 M HCl
HC2H3O2(aq) + NaOH(aq) → H2O(l) + NaC2H3O2(aq)
a. 16.58 × 10 −3 L soln ×
0.5062 mol NaOH 1 mol HC 2 H 3O 2
×
L soln
mol NaOH
= 8.393 × 10 −3 mol HC2H3O2
Concentration of HC2H3O2(aq) =
8.393 × 10 −3 mol
= 0.8393 M
0.01000 L
b. If we have 1.000 L of solution: Total mass = 1000. mL ×
Mass of HC2H3O2 = 0.8393 mol ×
Mass % acetic acid =
75.
1.006 g
= 1006 g solution
mL
60.05 g
= 50.40 g HC2H3O2
mol
50.40 g
× 100 = 5.010%
1006 g
2 HNO3(aq) + Ca(OH)2(aq) → 2 H2O(l) + Ca(NO3)2(aq)
35.00 × 10 −3 L HNO3 ×
0.0500 mol HNO3 1 mol Ca (OH) 2
1 L Ca (OH) 2
×
×
L HNO3
2 mol HNO3
0.0200 mol Ca (OH) 2
= 0.0438 L = 43.8 mL Ca(OH)2
76.
Strong bases contain the hydroxide ion (OH−). The reaction that occurs is H+ + OH− → H2O.
0.0120 L ×
0.150 mol H + 1 mol OH −
= 1.80 × 10 −3 mol OH−
×
L
mol H +
The 30.0 mL of the unknown strong base contains 1.80 × 10 −3 mol OH− .
1.80 × 10 −3 mol OH −
= 0.0600 M OH−
0.0300 L
The unknown base concentration is one-half the concentration of OH− ions produced from
the base, so the base must contain 2 OH− in each formula unit. The three soluble strong bases
that have two OH− ions in the formula are Ca(OH)2, Sr(OH)2, and Ba(OH)2. These are all
possible identities for the strong base.
CHAPTER 4
77.
SOLUTION STOICHIOMETRY
115
KHP is a monoprotic acid: NaOH(aq) + KHP(aq) →H2O(l) + NaKP(aq)
Mass KHP = 0.02046 L NaOH ×
0.1000 mol NaOH 1 mol KHP
204.22 g KHP
×
×
L NaOH
mol NaOH
mol KHP
= 0.4178 g KHP
78.
Because KHP is a monoprotic acid, the reaction is (KHP is an abbreviation for potassium
hydrogen phthalate):
NaOH(aq) + KHP(aq) → NaKP(aq) + H2O(l)
0.1082 g KHP ×
1 mol KHP
1 mol NaOH
×
= 5.298 × 10−4 mol NaOH
204.22 g KHP
mol KHP
There are 5.298 × 10−4 mol of sodium hydroxide in 34.67 mL of solution. Therefore, the
concentration of sodium hydroxide is:
5.298 × 10 −4 mol
= 1.528 × 10−2 M NaOH
−3
34.67 × 10 L
Oxidation-Reduction Reactions
79.
Apply the rules in Table 4.2.
a. KMnO4 is composed of K+ and MnO4− ions. Assign oxygen an oxidation state of −2,
which gives manganese a +7 oxidation state because the sum of oxidation states for all
atoms in MnO4− must equal the 1− charge on MnO4−. K, +1; O, −2; Mn, +7.
b. Assign O a −2 oxidation state, which gives nickel a +4 oxidation state. Ni, +4; O, −2.
c. Na4Fe(OH)6 is composed of Na+ cations and Fe(OH)64−anions. Fe(OH)64− is composed of
an iron cation and 6 OH− anions. For an overall anion charge of 4−, iron must have a +2
oxidation state. As is usually the case in compounds, assign O a −2 oxidation state and H
a +1 oxidation state. Na, +1; Fe, +2; O, −2; H, +1.
d. (NH4)2HPO4 is made of NH4+ cations and HPO42− anions. Assign +1 as the oxidation
state of H and −2 as the oxidation state of O. In NH4+, x + 4(+1) = +1, x = −3 = oxidation
state of N. In HPO42−, +1 + y + 4(−2) = −2, y = +5 = oxidation state of P.
e. O, −2; P, +3
f.
O, −2; 3x + 4(−2) = 0, x = +8/3 = oxidation state of Fe; this is the average oxidation state
of the three iron ions in Fe3O4. In the actual formula unit, there are two Fe3+ ions and one
Fe2+ ion.
g. O, −2; F, −1; Xe, +6
i.
O, −2; C, +2
h. F, −1; S, +4
j.
H, +1; O, −2; C, 0
116
80.
CHAPTER 4
SOLUTION STOICHIOMETRY
a. UO22+: O, −2; for U, x + 2(−2) = +2, x = +6
b. As2O3: O, −2; for As, 2(x) + 3(−2) = 0, x = +3
c. NaBiO3: Na, +1; O, −2; for Bi, +1 + x + 3(−2) = 0, x = +5
d. As4: As, 0
e. HAsO2: Assign H = +1 and O = −2; for As, +1 + x + 2(−2) = 0, x = +3
f.
Mg2P2O7: Composed of Mg2+ ions and P2O74− ions. Mg, +2; O, −2; P, +5
g. Na2S2O3: Composed of Na+ ions and S2O32− ions. Na, +1; O, −2; S, +2
h. Hg2Cl2: Hg, +1; Cl, −1
i.
81.
82.
Ca(NO3)2: Composed of Ca2+ ions and NO3− ions.
Ca, +2; O, −2; N, +5
a. −3
b. −3
c. 2(x) + 4(+1) = 0, x = −2
d. +2
e. +1
f.
+4
g. +3
h. +5
i.
0
a. SrCr2O7: Composed of Sr2+ and Cr2O72− ions. Sr, +2; O, −2; Cr, 2x + 7(−2) = −2, x = +6
b. Cu, +2; Cl, −1
c. O, 0;
d. H, +1; O, −1
e. Mg2+ and CO32− ions present. Mg, +2; O, −2; C, +4;
f.
g. Pb2+ and SO32− ions present. Pb, +2; O, −2; S, +4;
h. O, −2; Pb, +4
Ag, 0
i.
Na+ and C2O42− ions present. Na, +1; O, −2; C, 2x + 4(−2) = −2, x = +3
j.
O, −2; C, +4
k. Ammonium ion has a 1+ charge (NH4+), and sulfate ion has a 2− charge (SO42−).
Therefore, the oxidation state of cerium must be +4 (Ce4+). H, +1; N, −3; O, −2; S, +6
l.
83.
O, −2; Cr, +3
To determine if the reaction is an oxidation-reduction reaction, assign oxidation states. If the
oxidation states change for some elements, then the reaction is a redox reaction. If the
oxidation states do not change, then the reaction is not a redox reaction. In redox reactions,
the species oxidized (called the reducing agent) shows an increase in oxidation states, and the
species reduced (called the oxidizing agent) shows a decrease in oxidation states.
CHAPTER 4
SOLUTION STOICHIOMETRY
Redox?
117
Oxidizing
Agent
Reducing
Agent
Substance
Oxidized
Substance
Reduced
a. Yes
Ag+
Cu
Cu
Ag+
b. No
−
−
−
−
c. No
−
−
−
−
d. Yes
SiCl4
Mg
Mg
e. No
−
−
−
SiCl4 (Si)
−
In b, c, and e, no oxidation numbers change.
84.
The species oxidized shows an increase in oxidation states and is called the reducing agent.
The species reduced shows a decrease in oxidation states and is called the oxidizing agent.
The pertinent oxidation states are listed by the substance oxidized and the substance reduced.
Redox?
Oxidizing
Agent
Reducing
Agent
Substance
Oxidized
Substance
Reduced
a. Yes
H2O
CH4
CH4 (C, −4 → +2)
H2O (H, +1 → 0)
b
AgNO3
Cu
Cu (0 → +2)
AgNO3 (Ag, +1 → 0)
HCl
Zn
Zn (0 → +2)
HCl (H, +1 → 0)
Yes
c. Yes
d. No; there is no change in any of the oxidation numbers.
85.
Each sodium atom goes from the 0 oxidation state in Na to the +1 oxidation state in NaF.
Each Na atom loses one electron. Each fluorine atom goes from the 0 oxidation state in F2 to
the −1 state in NaF. In order to match electrons gained by fluorine with electrons lost by
sodium, 1 F atom is needed for every Na atom in the balanced equation. Because F2 contains
two fluorine atoms, two sodium atoms will be needed to balance the electrons. The following
balanced equation makes sense from an atom standpoint but also makes sense from an
electron standpoint.
2 Na(s) + F2(g) → 2 NaF(s)
86.
Each oxygen atom goes from the 0 oxidation state in O2 to the −2 oxidation state in MgO.
Each magnesium atom goes from the 0 oxidation state in Mg to the +2 oxidation state in
MgO. To match electron gain with electron loss, 1 atom of O is needed for each atom of Mg
in the balanced equation. Because two oxygen atoms are in each O2 molecule, we will need
two Mg atoms for every O2 molecule. The balanced equation below balances atoms but also
balances electrons, which must always be the case in any correctly balanced equation.
2 Mg(s) + O2(g) → 2 MgO(s)
87.
a. The first step is to assign oxidation states to all atoms (see numbers above the atoms).
−3 +1
0
+4 −2
+1 −2
C2H6 + O2 → CO2 + H2O
118
CHAPTER 4
SOLUTION STOICHIOMETRY
Each carbon atom changes from −3 to +4, an increase of 7. Each oxygen atom changes
from 0 to −2, a decrease of 2. We need 7/2 O atoms for every C atom in order to balance
electron gain with electron loss.
C2H6 + 7/2 O2 → CO2 + H2O
Balancing the remainder of the equation by inspection:
C2H6(g) + 7/2 O2(g) → 2 CO2(g) + 3 H2O(g)
or
2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(g)
b. The oxidation state of magnesium changes from 0 to +2, an increase of 2. The oxidation
state of hydrogen changes from +1 to 0, a decrease of 1. We need 2 H atoms for every
Mg atom in order to balance the electrons transferred. The balanced equation is:
Mg(s) + 2 HCl(aq) → Mg2+(aq) + 2 Cl−(aq) + H2(g)
c. The oxidation state of nickel increases by 2 (0 to +2), and the oxidation state of cobalt
decreases by 1 (+3 to +2). We need 2 Co3+ ions for every Ni atom in order to balance
electron gain with electron loss. The balanced equation is:
Ni(s) + 2 Co3+(aq) → Ni2+(aq) + 2 Co2+(aq)
d. The equation is balanced (mass and charge balanced). Each hydrogen atom gains one
electron (+1 → 0), and each zinc atom loses two electrons (0 → +2). We need 2 H atoms
for every Zn atom in order to balance the electrons transferred. This is the ratio in the
given equation:
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
88.
a. The first step is to assign oxidation states to all atoms (see numbers above the atoms).
0
0
−1
+3
Cl2 + Al → Al3+ + Cl−
Each aluminum atom changes in oxidation state from 0 to +3, an increase of 3. Each
chlorine atom changes from 0 to −1, a decrease of 1. We need 3 Cl atoms for every Al
atom in the balanced equation in order to balance electron gain with electron loss.
3/2 Cl2 + Al → Al3+ + 3 Cl−
For whole numbers, multiply through by two. The balanced equation is:
3 Cl2(g) + 2 Al(s) → 2Al3+(aq) + 6 Cl−(aq)
0
+1 −2
0
+2 −2 +1
b. O2 + H2O + Pb → Pb(OH)2
From the oxidation states written above the elements, lead is oxidized, and oxygen in O2
is reduced. Each lead atom changes from 0 to +2, an increase of 2, and each O atom in
CHAPTER 4
SOLUTION STOICHIOMETRY
119
O2 changes from 0 to −2, a decrease of 2. We need 1 Pb atom for each O atom in O2 to
balance the electrons transferred. Balancing the electrons:
O2 + H2O + 2 Pb → 2 Pb(OH)2
The last step is to balance the rest of the equation by inspection. In this reaction, when
the H atoms become balanced, the entire equation is balanced. The balanced overall
equation is:
O2(g) + 2 H2O(l) + 2 Pb(s) → 2 Pb(OH)2(s)
+7 −2
+1
+2
+2
+3
+1 −2
c. H+ + MnO4− + Fe2+ → Mn2+ + Fe3+ + H2O
From the oxidation states written above each element, manganese is reduced (goes from
+7 to +3), and Fe is oxidized (goes from +2 to +3). In order to balance the electrons
transferred, we need 5 Fe atoms for every Mn atom. Balancing the electrons gives:
H+ + MnO4− + 5 Fe2+ → Mn2+ + 5 Fe3+ + H2O
Balancing the O atoms, then the H atoms by inspection, leads to the following overall
balanced equation.
8 H+(aq) + MnO4−(aq) + 5 Fe2+(aq) → Mn2+(aq) + 5 Fe3+(aq) + 4 H2O(l)
Additional Exercises
89.
Desired uncertainty is 1% of 0.02, or ±0.0002. So we want the solution to be 0.0200 ±
0.0002 M, or the concentration should be between 0.0198 and 0.0202 M. We should use a 1L volumetric flask to make the solution. They are good to ±0.1%. We want to weigh out
between 0.0198 mol and 0.0202 mol of KIO3.
Molar mass of KIO3 = 39.10 + 126.9 + 3(16.00) = 214.0 g/mol
0.0198 mol ×
214.0 g
214.0 g
= 4.237 g; 0.0202 mol ×
= 4.323 g (carrying extra sig. figs.)
mol
mol
We should weigh out between 4.24 and 4.32 g of KIO3. We should weigh it to the nearest
milligram, or nearest 0.1 mg. Dissolve the KIO3 in water, and dilute (with mixing along the
way) to the mark in a 1-L volumetric flask. This will produce a solution whose concentration
is within the limits and is known to at least the fourth decimal place.
90.
Solution A:
6 molecules 1.5 molecules
4 molecules
=
; solution B:
1.0 L
4.0 L
1.0 L
Solution C:
4 molecules 2 molecules
6 molecules 3 molecules
=
=
; solution D:
2.0 L
1.0 L
2.0 L
1.0 L
120
CHAPTER 4
SOLUTION STOICHIOMETRY
Solution A has the most molecules per unit volume so solution A is most concentrated. This
is followed by solution D, then solution C. Solution B has the fewest molecules per unit
volume, so solution B is least concentrated.
91.
32.0 g C12H22O11 ×
1 mol C12 H 22 O11
= 0.0935 mol C12H22O11 added to blood
342.30 g
The blood sugar level would increase by:
0.0935 mol C12 H 22 O11
= 0.019 mol/L
5.0 L
92.
Mol CaCl2 present = 0.230 L CaCl2 ×
0.275 mol CaCl 2
= 6.33 × 10 −2 mol CaCl2
L CaCl 2
The volume of CaCl2 solution after evaporation is:
6.33 × 10 −2 mol CaCl2 ×
1 L CaCl 2
= 5.75 × 10 −2 L = 57.5 mL CaCl2
1.10 mol CaCl 2
Volume H2O evaporated = 230. mL − 57.5 mL = 173 mL H2O evaporated
93.
There are other possible correct choices for most of the following answers. We have listed
only three possible reactants in each case.
a. AgNO3, Pb(NO3)2, and Hg2(NO3)2 would form precipitates with the Cl− ion.
Ag+(aq) + Cl−(aq) → AgCl(s); Pb2+(aq) + 2 Cl−(aq) → PbCl2(s)
Hg22+(aq) + 2 Cl−(aq) →Hg2Cl2(s)
b. Na2SO4, Na2CO3, and Na3PO4 would form precipitates with the Ca2+ ion.
Ca2+(aq) + SO42−(aq) → CaSO4(s); Ca2+(aq) + CO32−(aq) → CaCO3(s)
3 Ca2+(aq) + 2 PO43−(aq) → Ca3(PO4)2(s)
c. NaOH, Na2S, and Na2CO3 would form precipitates with the Fe3+ ion.
Fe3+(aq) + 3 OH−(aq) → Fe(OH)3(s); 2 Fe3+(aq) + 3 S2−(aq) → Fe2S3(s)
2 Fe3+(aq) + 3 CO32−(aq) → Fe2(CO3)3(s)
d. BaCl2, Pb(NO3)2, and Ca(NO3)2 would form precipitates with the SO42− ion.
Ba2+(aq) + SO42−(aq) → BaSO4(s); Pb2+(aq) + SO42−(aq) → PbSO4(s)
Ca2+(aq) + SO42−(aq) → CaSO4(s)
e. Na2SO4, NaCl, and NaI would form precipitates with the Hg22+ ion.
Hg22+(aq) + SO42−(aq) → Hg2SO4(s); Hg22+(aq) + 2 Cl−(aq) → Hg2Cl2(s)
Hg22+(aq) + 2 I−(aq) → Hg2I2(s)
CHAPTER 4
f.
SOLUTION STOICHIOMETRY
121
NaBr, Na2CrO4, and Na3PO4 would form precipitates with the Ag+ ion.
Ag+(aq) + Br-(aq) → AgBr(s); 2 Ag+(aq) + CrO42−(aq) →Ag2CrO4(s)
3 Ag+(aq) + PO43−(aq) → Ag3PO4(s)
94.
a. MgCl2(aq) + 2 AgNO3(aq) → 2 AgCl(s) + Mg(NO3)2(aq)
1 mol MgCl 2
95.21 g
1 mol AgCl
×
×
= 0.213 g MgCl2
mol MgCl 2
2 mol AgCl
143.4 g AgCl
0.641 g AgCl ×
0.213 g MgCl 2
× 100 = 14.2% MgCl2
1.50 g mixture
b. 0.213 g MgCl2 ×
2 mol AgNO3
1 mol MgCl 2
1L
1000 mL
×
×
×
95.21 g
mol MgCl 2
0.500 mol AgNO3
1L
= 8.95 mL AgNO3
95.
XCl2(aq) + 2 AgNO3(aq) → 2 AgCl(s) + X(NO3)2(aq)
1.38 g AgCl ×
1 mol AgCl 1 mol XCl 2
×
= 4.81 × 10 −3 mol XCl2
143.4 g
2 mol AgCl
1.00 g XCl 2
4.91 × 10 −3 mol XCl 2
= 208 g/mol; x + 2(35.45) = 208, x = 137 g/mol
From the periodic table, the metal X is barium (Ba).
96.
From the periodic table, use aluminum in the formulas to convert from mass of Al(OH)3 to
mass of Al2(SO4)3 in the mixture.
0.107 g Al(OH)3 ×
1 mol Al2 (SO 4 ) 3
1 mol Al(OH) 3
1 mol Al3+
×
×
mol Al(OH) 3
78.00 g
2 mol Al3+
342.17 g Al2 (SO 4 ) 3
×
= 0.235 g Al2(SO4)3
mol Al2 (SO 4 ) 3
Mass % Al2(SO4)3 =
97.
0.235 g
× 100 = 16.2%
1.45 g
All the Tl in TlI came from Tl in Tl2SO4. The conversion from TlI to Tl2SO4 uses the molar
masses and formulas of each compound.
0.1824 g TlI ×
504.9 g Tl 2SO 4
204.4 g Tl
×
= 0.1390 g Tl2SO4
331.3 g TlI
408.8 g Tl
Mass % Tl2SO4 =
0.1390 g Tl 2SO 4
× 100 = 1.465% Tl2SO4
9.486 g pesticide
122
98.
CHAPTER 4
SOLUTION STOICHIOMETRY
a. Fe3+(aq) + 3 OH−(aq) → Fe(OH)3(s)
Fe(OH)3: 55.85 + 3(16.00) + 3(1.008) = 106.87 g/mol
0.107 g Fe(OH)3 ×
55.85 g Fe
= 0.0559 g Fe
106.87 g Fe(OH) 3
b. Fe(NO3)3: 55.85 + 3(14.01) + 9(16.00) = 241.86 g/mol
0.0559 g Fe ×
241.86 g Fe( NO 3 ) 3
= 0.242 g Fe(NO3)3
55.85 g Fe
c. Mass % Fe(NO3)3 =
99.
0.242 g
× 100 = 53.1%
0.456 g
With the ions present, the only possible precipitate is Cr(OH)3.
Cr(NO3)3(aq) + 3 NaOH(aq) → Cr(OH)3(s) + 3 NaNO3(aq)
Mol NaOH used = 2.06 g Cr(OH)3 ×
to form precipitate
1 mol Cr (OH) 3
3 mol NaOH
×
= 6.00 × 10−2 mol
mol Cr (OH) 3
103.02 g
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
Mol NaOH used = 0.1000 L ×
to react with HCl
MNaOH =
100.
0.400 mol HCl 1 mol NaOH
×
= 4.00 × 10−2 mol
L
mol HCl
total mol NaOH
6.00 × 10 −2 mol + 4.00 × 10 −2 mol
=
= 2.00 M NaOH
0.0500 L
volume
a. MgO(s) + 2 HCl(aq) → MgCl2(aq) + H2O(l)
Mg(OH)2(s) + 2 HCl(aq) → MgCl2(aq) + 2 H2O(l)
Al(OH)3(s) + 3 HCl(aq) → AlCl3(aq) + 3 H2O(l)
b. Let's calculate the number of moles of HCl neutralized per gram of substance. We can
get these directly from the balanced equations and the molar masses of the substances.
1 mol MgO
2 mol HCl
×
mol MgO
40.31 g MgO
2 mol HCl
mol Mg (OH) 2
×
=
4.962 × 10 −2 mol HCl
g MgO
1 mol Mg (OH) 2
58.33 g Mg (OH) 2
=
3.429 × 10 −2 mol HCl
g Mg (OH) 2
CHAPTER 4
SOLUTION STOICHIOMETRY
3 mol HCl
mol Al(OH) 3
×
1 mol Al(OH) 3
78.00 g Al(OH) 3
=
123
3.846 × 10 −2 mol HCl
g Al(OH) 3
Therefore, 1 gram of magnesium oxide would neutralize the most 0.10 M HCl.
101.
Using HA as an abbreviation for the monoprotic acid acetylsalicylic acid:
HA(aq) + NaOH(aq) → H2O(l) + NaA(aq)
Mol HA = 0.03517 L NaOH ×
0.5065 mol NaOH
1 mol HA
×
= 1.781 × 10−2 mol HA
L NaOH
mol NaOH
Fom the problem, 3.210 g HA was reacted, so:
molar mass =
102.
1.781 × 10 − 2 mol HA
= 180.2 g/mol
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
3.00 g Mg ×
103.
3.210 g HA
2 mol HCl
1L
1 mol Mg
×
×
= 0.049 L = 49 mL HCl
24.31 g Mg
mol Mg
5.0 mol HCl
Let HA = unknown monoprotic acid; HA(aq) + NaOH(aq) → NaA(aq) + H2O(l)
Mol HA present = 0.0250 L ×
0.500 mol NaOH
1 mol HA
×
L
1 mol NaOH
= 0.0125 mol HA
x g HA
2.20 g HA
=
, x = molar mass of HA = 176 g/mol
mol HA
0.0125 mol HA
Empirical formula weight ≈ 3(12) + 4(1) + 3(16) = 88 g/mol.
Because 176/88 = 2.0, the molecular formula is (C3H4O3)2 = C6H8O6.
104.
We get the empirical formula from the elemental analysis. Out of 100.00 g carminic acid,
there are:
53.66 g C ×
1 mol H
1 mol C
= 4.468 mol C; 4.09 g H ×
= 4.06 mol H
12.01 g C
1.008 g H
42.25 g O ×
1 mol O
= 2.641 mol O
16.00 g O
Dividing the moles by the smallest number gives:
4.468
= 1.692;
2.641
4.06
= 1.54
2.641
124
CHAPTER 4
SOLUTION STOICHIOMETRY
These numbers don’t give obvious mole ratios. Let’s determine the mol C to mol H ratio:
4.468
11
= 1.10 =
4.06
10
So let's try
4.468
4.06
4.06
2.641
= 0.406 as a common factor:
= 11.0;
= 10.0;
= 6.50
0.406
0.406
10
0.406
Therefore, C22H20O13 is the empirical formula.
We can get molar mass from the titration data. The balanced reaction is HA(aq) + OH−(aq)
→ H2O(l) + A−(aq), where HA is an abbreviation for carminic acid, an acid with one acidic
proton (H+).
0.0406 mol NaOH 1 mol carminic acid
×
L soln
mol NaOH
= 7.32 × 10−4 mol carminic acid
0.3602 g
492 g
=
Molar mass =
−4
mol
7.32 × 10 mol
18.02 × 10−3 L soln ×
The empirical formula mass of C22H20O13 ≈ 22(12) + 20(1) + 13(16) = 492 g.
Therefore, the molecular formula of carminic acid is also C22H20O13.
105.
0.104 g AgCl ×
1 mol AgCl
35.45 g Cl −
1 mol Cl −
= 2.57 × 10−2 g Cl−
×
×
143.4 g AgCl
mol AgCl
mol Cl −
All of the Cl− in the AgCl precipitate came from the chlorisondamine chloride compound in
the medication. So we need to calculate the quantity of C14H20Cl6N2 which contains 2.57 ×
10−2 g Cl−.
Molar mass of C14H20Cl6N2 = 14(12.01) + 20(1.008) + 6(35.45) + 2(14.01) = 429.02 g/mol
There are 6(35.45) = 212.70 g chlorine for every mole (429.02 g) of C14H20Cl6N2.
2.57 × 10−2 g Cl− ×
429.02 g C14 H 20 Cl 6 N 2
212.70 g Cl
Mass % chlorisondamine chloride =
106.
−
= 5.18 × 10−2 g C14H20Cl6N2
5.18 × 10 −2 g
× 100 = 4.05%
1.28 g
All the sulfur in BaSO4 came from the saccharin. The conversion from BaSO4 to saccharin
utilizes the molar masses of each compound.
0.5032 g BaSO4 ×
183.19 g C 7 H 5 NO3S
32.07 g S
×
= 0.3949 g C7H5NO3S
233.4 g BaSO 4
32.07 g S
Average mass
0.3949 g
3.949 × 10 −2 g
39.49 mg
=
=
=
Tablet
10 tablets
tablet
tablet
CHAPTER 4
SOLUTION STOICHIOMETRY
Average mass % =
107.
0.3949 g C 7 H 5 NO 3S
× 100 = 67.00% saccharin by mass
0.5894 g
Use the silver nitrate data to calculate the mol Cl− present, then use the formula of douglasite
(2KCl•FeCl2•2H2O) to convert from Cl− to douglasite (1 mole of douglasite contains 4 moles
of Cl−). The net ionic equation is Ag+ + Cl− → AgCl(s).
0.03720 L ×
0.1000 mol Ag + 1 mol Cl − 1 mol douglasite 311.88 g douglasite
×
×
×
L
mol
mol Ag +
4 mol Cl −
= 0.2900 g douglasite
Mass % douglasite =
108.
125
0.2900 g
× 100 = 63.74%
0.4550 g
a. Al(s) + 3 HCl(aq) → AlCl3(aq) + 3/2 H2(g) or 2 Al(s) + 6 HCl(aq) → 2 AlCl3(aq) +
3 H2(g)
Hydrogen is reduced (goes from the +1 oxidation state to the 0 oxidation state), and
aluminum Al is oxidized (0 → +3).
b. Balancing S is most complicated since sulfur is in both products. Balance C and H first;
then worry about S.
CH4(g) + 4 S(s) → CS2(l) + 2 H2S(g)
Sulfur is reduced (0 → −2), and carbon is oxidized (−4 → +4).
c. Balance C and H first; then balance O.
C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(l)
Oxygen is reduced (0 → −2), and carbon is oxidized (−8/3 → +4).
d. Although this reaction is mass balanced, it is not charge balanced. We need 2 moles of
silver on each side to balance the charge.
Cu(s) + 2 Ag+(aq) → 2 Ag(s) + Cu2+(aq)
Silver is reduced (+1 → 0), and copper is oxidized (0 → +2).
109.
Cr2O72−: 2(x) + 7(−2) = −2, x = +6
C2H5OH (C2H6O): 2(y) + 6(+1) + (−2) = 0, y = −2
CO2: z + 2(−2) = 0, z = +4
Each chromium atom goes from the oxidation state of +6 in Cr2O72− to +3 in Cr3+. Each
chromium atom gains three electrons; chromium is the species reduced. Each carbon atom
goes from the oxidation state of −2 in C2H5OH to +4 in CO2. Each carbon atom loses six
126
CHAPTER 4
SOLUTION STOICHIOMETRY
electrons; carbon is the species oxidized. From the balanced equation, we have four
chromium atoms and two carbon atoms. With each chromium atom gaining three electrons, a
total of 4(3) = 12 electrons are transferred in the balanced reaction. This is confirmed from
the 2 carbon atoms in the balanced equation, where each carbon atom loses six electrons [2(6)
= 12 electrons transferred].
ChemWork Problems
The answers to the problems 110-119 (or a variation to these problems) are found in OWL. These
problems are also assignable in OWL.
Challenge Problems
120.
Let x = mass of NaCl, and let y = mass K2SO4. So x + y = 10.00.
Two reactions occur: Pb2+(aq) + 2 Cl−(aq) → PbCl2(s) and
Pb2+(aq) + SO42−(aq) → PbSO4(s)
Molar mass of NaCl = 58.44 g/mol; molar mass of K2SO4 = 174.27 g/mol; molar mass of
PbCl2 = 278.1 g/mol; molar mass of PbSO4 = 303.3 g/mol
y
x
= moles NaCl;
= moles K2SO4
174.27
58.44
mass of PbCl2 + mass PbSO4 = total mass of solid
x
y
(1/2)(278.1) +
(303.3) = 21.75
58.44
174.27
We have two equations:
(2.379)x + (1.740)y = 21.75 and x + y = 10.00. Solving:
x = 6.81 g NaCl;
121.
6.81 g NaCl
× 100 = 68.1% NaCl
10.00 g mixture
a. 5.0 ppb Hg in water =
5.0 ng Hg
5.0 × 10 −9 g Hg
=
g soln
mL soln
5.0 × 10 −9 g Hg
1 mol Hg
1000 mL
= 2.5 × 10 −8 M Hg
×
×
mL
200.6 g Hg
L
b.
1.0 × 10 −9 g CHCl3
1 mol CHCl3
1000 mL
= 8.4 × 10 −9 M CHCl3
×
×
mL
119.37 g CHCl3
L
c. 10.0 ppm As =
10.0 μg As 10.0 × 10 −6 g As
=
g soln
mL soln
CHAPTER 4
SOLUTION STOICHIOMETRY
127
10.0 × 10 −6 g As
1 mol As
1000 mL
= 1.33 × 10 −4 M As
×
×
mL
74.92 g As
L
d.
122.
0.10 × 10 −6 g DDT
1 mol DDT
1000 mL
= 2.8 × 10 −7 M DDT
×
×
mL
354.46 g DDT
L
We want 100.0 mL of each standard. To make the 100. ppm standard:
100. μg Cu
× 100.0 mL solution = 1.00 × 104 µg Cu needed
mL
1.00 × 104 µg Cu ×
1 mL stock
= 10.0 mL of stock solution
1000.0 μg Cu
Therefore, to make 100.0 mL of 100. ppm solution, transfer 10.0 mL of the 1000.0 ppm stock
solution to a 100-mL volumetric flask, and dilute to the mark.
Similarly:
75.0 ppm standard, dilute 7.50 mL of the 1000.0 ppm stock to 100.0 mL.
50.0 ppm standard, dilute 5.00 mL of the 1000.0 ppm stock to 100.0 mL.
25.0 ppm standard, dilute 2.50 mL of the 1000.0 ppm stock to 100.0 mL.
10.0 ppm standard, dilute 1.00 mL of the 1000.0 ppm stock to 100.0 mL.
123.
a. 0.308 g AgCl ×
35.45 g Cl
0.0761 g
= 0.0761 g Cl; % Cl =
× 100 = 29.7% Cl
143.4 g AgCl
0.256 g
Cobalt(III) oxide, Co2O3: 2(58.93) + 3(16.00) = 165.86 g/mol
0.145 g Co2O3 ×
117.86 g Co
0.103 g
= 0.103 g Co; % Co =
× 100 = 24.8% Co
0.416 g
165.86 g Co 2 O 3
The remainder, 100.0 − (29.7 + 24.8) = 45.5%, is water.
Assuming 100.0 g of compound:
45.5 g H2O ×
2.016 g H
5.09 g H
= 5.09 g H; % H =
× 100 = 5.09% H
18.02 g H 2 O
100.0 g compound
45.5 g H2O ×
16.00 g O
40.4 g O
= 40.4 g O; % O =
× 100 = 40.4% O
18.02 g H 2 O
100.0 g compound
The mass percent composition is 24.8% Co, 29.7% Cl, 5.09% H, and 40.4% O.
b. Out of 100.0 g of compound, there are:
128
CHAPTER 4
24.8 g Co Η
5.09 g H ×
SOLUTION STOICHIOMETRY
1 mol
1 mol
= 0.421 mol Co; 29.7 g Cl ×
= 0.838 mol Cl
58.93 g Co
35.45 g Cl
1 mol
1 mol
= 5.05 mol H; 40.4 g O ×
= 2.53 mol O
1.008 g H
16.00 g O
Dividing all results by 0.421, we get CoCl2•6H2O for the empirical formula, which is also
the actual formula given the information in the problem. The •6H2O represent six waters
of hydration in the chemical formula.
c. CoCl2•6H2O(aq) + 2 AgNO3(aq) → 2 AgCl(s) + Co(NO3)2(aq) + 6 H2O(l)
CoCl2•6H2O(aq) + 2 NaOH(aq) → Co(OH)2(s) + 2 NaCl(aq) + 6 H2O(l)
Co(OH)2 → Co2O3
This is an oxidation-reduction reaction. Thus we also need to
include an oxidizing agent. The obvious choice is O2.
4 Co(OH)2(s) + O2(g) → 2 Co2O3(s) + 4 H2O(l)
124.
a. C12H10-nCln + n Ag+ → n AgCl; molar mass of AgCl = 143.4 g/mol
Molar mass of PCB = 12(12.01) + (10 − n)(1.008) + n(35.45) = 154.20 + (34.44)n
Because n mol AgCl is produced for every 1 mol PCB reacted, n(143.4) g of AgCl will
be produced for every [154.20 + (34.44)n] g of PCB reacted.
Mass of AgCl
(143.4)n
=
or massAgCl[154.20 + (34.44)n] = massPCB(143.4)n
Mass of PCB
154.20 + (34.44)n
b. 0.4971[154.20 + (34.44)n] = 0.1947(143.4)n, 76.65 + (17.12)n = (27.92)n
76.65 = (10.80)n, n = 7.097
125.
Zn(s) + 2 AgNO2(aq) → 2 Ag(s) + Zn(NO2)2(aq)
Let x = mass of Ag and y = mass of Zn after the reaction has stopped. Then x + y = 29.0 g.
Because the moles of Ag produced will equal two times the moles of Zn reacted:
(19.0 − y) g Zn ×
1 mol Zn
2 mol Ag
1 mol Ag
×
= x g Ag ×
65.38 g Zn 1 mol Zn
107.9 g Ag
Simplifying:
3.059 × 10−2(19.0 − y) = (9.268 × 10−3)x
Substituting x = 29.0 − y into the equation gives:
3.059 × 10−2(19.0 − y) = 9.268 × 10−3(29.0 − y)
CHAPTER 4
SOLUTION STOICHIOMETRY
129
Solving:
0.581 − (3.059 × 10−2)y = 0.269 − (9.268 × 10−3)y, (2.132 × 10−2)y = 0.312, y = 14.6 g Zn
14.6 g Zn is present, and 29.0 − 14.6 = 14.4 g Ag is also present after the reaction is stopped.
126.
Ag+(aq) + Cl−(aq) → AgCl(s); let x = mol NaCl and y = mol KCl.
(22.90 × 10−3 L) × 0.1000 mol/L = 2.290 × 10−3 mol Ag+ = 2.290 × 10−3 mol Cl− total
x + y = 2.290 × 10−3 mol Cl−, x = 2.290 × 10−3 − y
Because the molar mass of NaCl is 58.44 g/mol and the molar mass of KCl is 74.55 g/mol:
(58.44)x + (74.55)y = 0.1586 g
58.44(2.290 × 10−3 − y) + (74.55)y = 0.1586, (16.11)y = 0.0248, y = 1.54 × 10−3 mol KCl
Mass % KCl =
1.54 × 10 −3 mol × 74.55 g / mol
× 100 = 72.4% KCl
0.1586 g
% NaCl = 100.0 − 72.4 = 27.6% NaCl
2−
127.
0.298 g BaSO4 ×
96.07 g SO 4
0.123 g SO 4
= 0.123 g SO42−; % sulfate =
233.4 g BaSO 4
0.205 g
2−
= 60.0%
Assume we have 100.0 g of the mixture of Na2SO4 and K2SO4. There are:
60.0 g SO42− ×
1 mol
= 0.625 mol SO42−
96.07 g
There must be 2 × 0.625 = 1.25 mol of 1+ cations to balance the 2− charge of SO42−.
Let x = number of moles of K+ and y = number of moles of Na+; then x + y = 1.25.
The total mass of Na+ and K+ must be 40.0 g in the assumed 100.0 g of mixture. Setting up
an equation:
x mol K+ ×
39.10 g
22.99 g
+ y mol Na+ ×
= 40.0 g
mol
mol
So we have two equations with two unknowns: x + y = 1.25 and (39.10)x + (22.99)y = 40.0
x = 1.25 − y, so 39.10(1.25 − y) + (22.99)y = 40.0
48.9 − (39.10)y + (22.99)y = 40.0, − (16.11)y = −8.9
y = 0.55 mol Na+ and x = 1.25 − 0.55 = 0.70 mol K+
130
CHAPTER 4
SOLUTION STOICHIOMETRY
Therefore:
0.70 mol K+ ×
1 mol K 2 SO 4
2 mol K
+
= 0.35 mol K2SO4; 0.35 mol K2SO4 ×
174.27 g
mol
= 61 g K2SO4
We assumed 100.0 g; therefore, the mixture is 61% K2SO4 and 39% Na2SO4.
128.
a. Let x = mass of Mg, so 10.00 − x = mass of Zn. Ag+(aq) + Cl−(aq) → AgCl(s).
From the given balanced equations, there is a 2 : 1 mole ratio between mol Mg and mol
Cl−. The same is true for Zn. Because mol Ag+ = mol Cl− present, one can set up an
equation relating mol Cl− present to mol Ag+ added.
x g Mg ×
1 mol Mg
2 mol Cl −
1 mol Zn
2 mol Cl −
×
+ (10.00 − x) g Zn ×
×
24.31 g Mg
mol Mg
65.38 g Zn
mol Zn
= 0.156 L ×
3.00 mol Ag + 1 mol Cl −
= 0.468 mol Cl−
×
+
L
mol Ag
20.00 − 2 x
2(10.00 − x)
2x
 2x

= 0.468, 24.31 × 65.38 
= 0.468 
+
+
65.38
65.38
24.31
 24.31

(130.8)x + 486.2 − (48.62)x = 743.8 (carrying 1 extra sig. fig.)
(82.2)x = 257.6, x = 3.13 g Mg;
b. 0.156 L ×
MHCl =
129.
% Mg =
3.13 g Mg
× 100 = 31.3% Mg
10.00 g mixture
3.00 mol Ag + 1 mol Cl −
= 0.468 mol Cl− = 0.468 mol HCl added
×
+
L
mol Ag
0.468 mol
= 6.00 M HCl
0.0780 L
Pb2+(aq) + 2 Cl−(aq) → PbCl2(s)
3.407 g PbCl2 ×
0.01225 mol
2.00 × 10
−3
L
1 mol PbCl 2
278.1 g PbCl 2
×
1 mol Pb 2+
= 0.01225 mol Pb2+
mol PbCl 2
= 6.13 M Pb2+ = 6.13 M Pb(NO3)2
This is also the Pb(NO3)2 concentration in the 80.0 mL of evaporated solution.
130.
Original concentration =
moles Pb( NO3 ) 2
0.0800 L × 6.13 mol/L
=
= 4.90 M Pb(NO3)2
0.1000 L
original volume
Mol CuSO4 = 87.7 mL ×
1L
0.500 mol
×
= 0.0439 mol
1000 mL
L
CHAPTER 4
SOLUTION STOICHIOMETRY
Mol Fe = 2.00 g ×
131
1 mol Fe
= 0.0358 mol
55.85 g
The two possible reactions are:
I.
CuSO4(aq) + Fe(s) → Cu(s) + FeSO4(aq)
II. 3 CuSO4(aq) + 2 Fe(s) → 3 Cu(s) + Fe2(SO4)3(aq)
If reaction I occurs, Fe is limiting, and we can produce:
0.0358 mol Fe ×
1 mol Cu 63.55 g Cu
×
= 2.28 g Cu
mol Fe
mol Cu
If reaction II occurs, CuSO4 is limiting, and we can produce:
0.0439 mol CuSO4 ×
3 mol Cu
63.55 g Cu
×
= 2.79 g Cu
3 mol CuSO 4
mol Cu
Assuming 100% yield, reaction I occurs because it fits the data best.
131.
0.2750 L × 0.300 mol/L = 0.0825 mol H+; let y = volume (L) delivered by Y and z
= volume (L) delivered by Z.
H+(aq) + OH−(aq) → H2O(l); y(0.150 mol/L) + z(0.250 mol/L) = 0.0825 mol H+
mol OH−
0.2750 L + y + z = 0.655 L, y + z = 0.380, z = 0.380 − y
y(0.150) + (0.380 − y)(0.250) = 0.0825, solving: y = 0.125 L, z = 0.255 L
Flow rate for Y =
132.
125 mL
255 mL
= 2.06 mL/min; flow rate for Z =
= 4.20 mL/min
60.65 min
60.65 min
a. H3PO4(aq) + 3 NaOH(aq) → 3 H2O(l) + Na3PO4(aq)
b. 3 H2SO4(aq) + 2 Al(OH)3(s) → 6 H2O(l) + Al2(SO4)3(aq)
c. H2Se(aq) + Ba(OH)2(aq) → 2 H2O(l) + BaSe(s)
d. H2C2O4(aq) + 2 NaOH(aq) →2 H2O(l) + Na2C2O4(aq)
133.
2 H3PO4(aq) + 3 Ba(OH)2(aq) → 6 H2O(l) + Ba3(PO4)2(s)
0.01420 L ×
0.141 mol H 3 PO 4
1 L Ba (OH) 2
3 mol Ba (OH) 2
= 0.0576 L
×
×
2 mol H 3 PO 4
0.0521 mol Ba (OH) 2
L
= 57.6 mL Ba(OH)2
132
134.
CHAPTER 4
35.08 mL NaOH ×
Molarity =
135.
SOLUTION STOICHIOMETRY
1 mol H 2SO 4
1L
2.12 mol NaOH
×
×
= 3.72 × 10 −2 mol H2SO4
1000 mL
L NaOH
2 mol NaOH
3.72 × 10 −2 mol
1000 mL
×
10.00 mL
L
= 3.72 M H2SO4
The pertinent equations are:
2 NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2 H2O(l)
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Amount of NaOH added = 0.0500 L ×
0.213 mol
= 1.07 × 10−2 mol NaOH
L
Amount of NaOH neutralized by HCl:
0.103 mol HCl 1 mol NaOH
×
= 1.36 × 10−3 mol NaOH
L HCl
mol HCl
0.01321 L HCl Η
The difference, 9.3 × 10−3 mol, is the amount of NaOH neutralized by the sulfuric acid.
9.3 × 10−3 mol NaOH ×
1 mol H 2SO 4
= 4.7 × 10−3 mol H2SO4
2 mol NaOH
Concentration of H2SO4 =
136.
4.7 × 10 −3 mol
= 4.7 × 10−2 M H2SO4
0.1000 L
Let H2A = formula for the unknown diprotic acid.
H2A(aq) + 2 NaOH(aq) → 2 H2O(l) + Na2A(aq)
Mol H2A = 0.1375 L ×
Molar mass of H2A =
137.
1 mol H 2 A
0.750 mol NaOH
×
L
2 mol NaOH
= 0.0516 mol
6.50 g
= 126 g/mol
0.0516 mol
Mol C6H8O7 = 0.250 g C6H8O7 ×
1 mol C 6 H 8 O 7
= 1.30 × 10 −3 mol C6H8O7
192.12 g C 6 H 8 O 7
Let HxA represent citric acid, where x is the number of acidic hydrogens. The balanced
neutralization reaction is:
HxA(aq) + x OH−(aq) → x H2O(l) + Ax−(aq)
Mol OH− reacted = 0.0372 L ×
0.105 mol OH −
= 3.91 × 10 −3 mol OH−
L
CHAPTER 4
x =
SOLUTION STOICHIOMETRY
3.91 × 10 −3 mol
mol OH −
=
mol citric acid
1.30 × 10 −3 mol
133
= 3.01
Therefore, the general acid formula for citric acid is H3A, meaning that citric acid has three
acidic hydrogens per citric acid molecule (citric acid is a triprotic acid).
138.
a. Flow rate = 5.00 × 104 L/s + 3.50 × 103 L/s = 5.35 × 104 L/s
b. CHCl =
3.50 × 103 (65.0)
= 4.25 ppm HCl
5.35 × 10 4
c. 1 ppm = 1 mg/kg H2O = 1 mg/L (assuming density = 1.00 g/mL)
8.00 h ×
60 min 60 s 1.80 × 10 4 L 4.25 mg HCl
1g
= 2.20 × 106 g HCl
×
×
×
×
h
min
s
L
1000 mg
2.20 × 106 g HCl ×
1 mol HCl
1 mol CaO 56.08 g Ca
×
×
= 1.69 × 106 g CaO
36.46 g HCl 2 mol HCl
mol CaO
d. The concentration of Ca2+ going into the second plant was:
5.00 × 10 4 (10.2)
= 9.53 ppm
5.35 × 10 4
The second plant used: 1.80 × 104 L/s × (8.00 × 60 × 60) s = 5.18 × 108 L of water.
1.69 × 106 g CaO ×
40.08 g Ca 2+
= 1.21 × 106 g Ca2+ was added to this water.
56.08 g CaO
C Ca 2+ (plant water) = 9.53 +
1.21 × 109 mg
= 9.53 + 2.34 = 11.87 ppm
5.18 × 108 L
Because 90.0% of this water is returned, (1.80 × 104) × 0.900 = 1.62 × 104 L/s of water
with 11.87 ppm Ca2+ is mixed with (5.35 − 1.80) × 104 = 3.55 × 104 L/s of water
containing 9.53 ppm Ca2+.
C Ca 2+ (final) =
139.
(1.62 × 10 4 L / s)(11.87 ppm) + (3.55 × 10 4 L / s)(9.53 ppm)
= 10.3 ppm
1.62 × 10 4 L / s + 3.55 × 10 4 L / s
Mol KHP used = 0.4016 g ×
1 mol
= 1.967 × 10 −3 mol KHP
204.22 g
Because 1 mole of NaOH reacts completely with 1 mole of KHP, the NaOH solution contains
1.967 × 10 −3 mol NaOH.
Molarity of NaOH =
1.967 × 10 −3 mol
25.06 × 10 −3 L
=
7.849 × 10 −2 mol
L
134
CHAPTER 4
Maximum molarity =
Minimum molarity =
1.967 × 10 −3 mol
25.01 × 10 −3 L
1.967 × 10 −3 mol
25.11 × 10 −3 L
SOLUTION STOICHIOMETRY
=
7.865 × 10 −2 mol
L
=
7.834 × 10 −2 mol
L
We can express this as 0.07849 ±0.00016 M. An alternative way is to express the molarity as
0.0785 ±0.0002 M. This second way shows the actual number of significant figures in the
molarity. The advantage of the first method is that it shows that we made all our individual
measurements to four significant figures.
Integrative Problems
140.
a. Assume 100.00 g of material.
42.23 g C ×
2.11 g B ×
1 mol C
1 mol F
= 3.516 mol C; 55.66 g F ×
= 2.929 mol F
12.01 g C
19.00 g F
1 mol B
= 0.195 mol B
10.81 g B
Dividing by the smallest number:
3.516
2.929
= 18.0;
= 15.0
0.195
0.195
The empirical formula is C18F15B.
b. 0.3470 L ×
0.01267 mol
= 4.396 × 10 −3 mol BARF
L
Molar mass of BARF =
2.251 g
= 512.1 g/mol
4.396 × 10 −3 mol
The empirical formula mass of BARF is 511.99 g. Therefore, the molecular formula is
the same as the empirical formula, C18F15B.
141.
3 (NH4)2CrO4(aq) + 2 Cr(NO2)3(aq) → 6 NH4NO2(aq) + Cr2(CrO4)3(s)
0.203 L × 0.307 mol ( NH 4 ) 2 CrO 4 × 1 mol Cr2 (CrO 4 ) 3 × 452.00 g Cr2 (CrO 4 ) 3
L
3 mol ( NH 4 ) 2 CrO 4
mol Cr2 (CrO 4 ) 3
= 9.39 g Cr2(CrO4)3
0.137 L × 0.269 mol Cr ( NO 2 ) 3 × 1 mol Cr2 (CrO 4 ) 3 × 452.00 g Cr2 (CrO 4 ) 3
L
2 mol Cr ( NO 2 ) 3
mol Cr2 (CrO 4 ) 3
= 8.33 g Cr2(CrO4)3
The Cr(NO2)3 reagent produces the smaller amount of product, so Cr(NO2)3 is limiting and
the theoretical yield of Cr2(CrO4)3 is 8.33 g.
CHAPTER 4
0.880 =
142.
SOLUTION STOICHIOMETRY
135
actual yield
, actual yield = (8.33 g)(0.880) = 7.33 g Cr2(CrO4)3 isolated
8.33 g
The oxidation states of the elements in the various ions are:
VO2+: O, −2; V, x + (−2) = + 2, x = +4
MnO4−: O, −2; Mn, x + 4(−2) = −1, x = +7
V(OH)4+: O, −2, H, +1; V, x + 4(−2) + 4(+1) = +1, x = +5
Mn2+: Mn, +2
Vanadium goes from the +4 oxidation state in VO2+ to the +5 oxidation state in V(OH)4+.
Manganese goes from the +7 oxidation state in MnO4− to the +2 oxidation state in Mn2+. We
need 5 V atoms for every Mn atom in order to balance the electrons transferred. Balancing
the electrons transferred, then balancing the rest by inspection gives:
MnO4−(aq) + 5 VO2+(aq) + 11 H2O(l) → 5 V(OH)4+(aq) + Mn2+(aq) + 22 H+(aq)
0.02645 L ×
0.581 =
143.
0.02250 mol MnO 4 −
5 mol VO 2+
1 mol V
50.94 g V
= 0.1516 g V
×
×
×
−
2
+
L
mol V
mol VO
mol MnO 4
0.1516 g V
, mass of ore sample = 0.1516/0.581 = 0.261 g
mass of ore sample
X2− contains 36 electrons, so X2− has 34 protons, which identifies X as selenium (Se). The
name of H2Se would be hydroselenic acid following the conventions described in Chapter 2.
H2Se(aq) + 2 OH−(aq) → Se2−(aq) + 2 H2O(l)
0.0356 L ×
0.175 mol OH − 1 mol H 2Se 80.98 g H 2Se
= 0.252 g H2Se
×
×
mol H 2Se
L
2 mol OH −
Marathon Problems
144.
Mol BaSO4 = 0.2327 g ×
1 mol
= 9.970 × 10 −4 mol BaSO4
233.4 g
The moles of the sulfate salt depend on the formula of the salt. The general equation is:
Mx(SO4)y(aq) + y Ba2+(aq) → y BaSO4(s) + x Mz+
Depending on the value of y, the mole ratio between the unknown sulfate salt and BaSO4
varies. For example, if Pat thinks the formula is TiSO4, the equation becomes:
TiSO4(aq) + Ba2+(aq) → BaSO4(s) + Ti2+(aq)
136
CHAPTER 4
SOLUTION STOICHIOMETRY
Because there is a 1 : 1 mole ratio between mol BaSO4 and mol TiSO4, you need 9.970 ×
10 −4 mol of TiSO4. Because 0.1472 g of salt was used, the compound would have a molar
mass of (assuming the TiSO4 formula):
0.1472 g/9.970 × 10 −4 mol = 147.6 g/mol
From atomic masses in the periodic table, the molar mass of TiSO4 is 143.95 g/mol. From
just these data, TiSO4 seems reasonable.
Chris thinks the salt is sodium sulfate, which would have the formula Na2SO4. The equation
is:
Na2SO4(aq) + Ba2+(aq) → BaSO4(s) + 2 Na+(aq)
As with TiSO4, there is a 1:1 mole ratio between mol BaSO4 and mol Na2SO4. For sodium
sulfate to be a reasonable choice, it must have a molar mass of about 147.6 g/mol. Using
atomic masses, the molar mass of Na2SO4 is 142.05 g/mol. Thus Na2SO4 is also reasonable.
Randy, who chose gallium, deduces that gallium should have a 3+ charge (because it is in
column 3A), and the formula of the sulfate would be Ga2(SO4)3. The equation would be:
Ga2(SO4)3(aq) + 3 Ba2+(aq) → 3 BaSO4(s) + 2 Ga3+(aq)
The calculated molar mass of Ga2(SO4)3 would be:
0.1472 g Ga 2 (SO 4 ) 3
3 mol BaSO 4
×
= 442.9 g/mol
−4
mol Ga 2 (SO 4 ) 3
9.970 × 10 mol BaSO 4
Using atomic masses, the molar mass of Ga2(SO4)3 is 427.65 g/mol. Thus Ga2(SO4)3 is also
reasonable.
Looking in references, sodium sulfate (Na2SO4) exists as a white solid with orthorhombic
crystals, whereas gallium sulfate [Ga2(SO4)3] is a white powder. Titanium sulfate exists as a
green powder, but its formula is Ti2(SO4)3. Because this has the same formula as gallium
sulfate, the calculated molar mass should be around 443 g/mol. However, the molar mass of
Ti2(SO4)3 is 383.97 g/mol. It is unlikely, then, that the salt is titanium sulfate.
To distinguish between Na2SO4 and Ga2(SO4)3, one could dissolve the sulfate salt in water
and add NaOH. Ga3+ would form a precipitate with the hydroxide, whereas Na2SO4 would
not. References confirm that gallium hydroxide is insoluble in water.
145.
a. Compound A = M(NO3)x; in 100.00 g of compd.: 8.246 g N ×
48.00 g O
= 28.25 g O
14.01 g N
Thus the mass of nitrate in the compound = 8.246 + 28.25 g = 36.50 g (if x = 1).
If x = 1: mass of M = 100.00 − 36.50 g = 63.50 g
Mol M = mol N =
8.246 g
= 0.5886 mol
14.01 g / mol
CHAPTER 4
SOLUTION STOICHIOMETRY
Molar mass of metal M =
137
63.50 g
= 107.9 g/mol (This is silver, Ag.)
0.5886 mol
If x = 2: mass of M = 100.00 − 2(36.50) = 27.00 g
Mol M = ½ mol N =
0.5886 mol
= 0.2943 mol
2
Molar mass of metal M =
27.00 g
= 91.74 g/mol
0.2943 mol
This is close to Zr, but Zr does not form stable 2+ ions in solution; it forms stable 4+
ions. Because we cannot have x = 3 or more nitrates (three nitrates would have a mass
greater than 100.00 g), compound A must be AgNO3.
Compound B: K2CrOx is the formula. This salt is composed of K+ and CrOx2− ions. Using
oxidation states, 6 + x(−2) = −2, x = 4. Compound B is K2CrO4 (potassium chromate).
b. The reaction is:
2 AgNO3(aq) + K2CrO4(aq) → Ag2CrO4(s) + 2 KNO3(aq)
The blood red precipitate is Ag2CrO4(s).
c. 331.8 g Ag2CrO4 formed; this is equal to the molar mass of Ag2CrO4, so 1 mole of
precipitate formed. From the balanced reaction, we need 2 mol AgNO3 to react with
1 mol K2CrO4 to produce 1 mol (331.8 g) of Ag2CrO4.
2.000 mol AgNO3 ×
169.9 g
= 339.8 g AgNO3
mol
1.000 mol K2CrO4 ×
194.2 g
= 194.2 g K2CrO4
mol
The problem says that we have equal masses of reactants. Our two choices are 339.8 g
AgNO3 + 339.8 g K2CrO4 or 194.2 g AgNO3 + 194.2 g K2CrO4. If we assume the 194.2-g
quantities are correct, then when 194.2 g K2CrO4 (1 mol) reacts, 339.8 g AgNO3 (2.0
mol) must be present to react with all the K2CrO4. We only have 194.2 g AgNO3 present;
this cannot be correct. Instead of K2CrO4 limiting, AgNO3 must be limiting, and we have
reacted 339.8 g AgNO3 and 339.8 g K2CrO4.
Solution A:
2.000 mol NO 3
2.000 mol Ag +
= 4.000 M Ag+ ;
0.5000 L
0.5000 L
Solution B: 339.8 g K2CrO4 ×
−
= 4.000 M NO3−
1 mol
= 1.750 mol K2CrO4
194.2 g
1.750 mol CrO 4
2 × 1.750 mol K +
= 7.000 M K+;
0.5000 L
0.5000 L
2−
= 3.500 M CrO42−
138
CHAPTER 4
SOLUTION STOICHIOMETRY
d. After the reaction, moles of K+ and moles of NO3− remain unchanged because they are
spectator ions. Because Ag+ is limiting, its concentration will be 0 M after precipitation is
complete. The following summarizes the changes that occur as the precipitate forms.
2 Ag+(aq)
Initial
2.000 mol
Change −2.000 mol
After rxn
0
+ CrO42−(aq)
1.750 mol
−1.000 mol
0.750 mol
→ Ag2CrO4(s)
0
+1.000 mol
1.000 mol
M K+ =
2 × 1.750 mol
2.000 mol
= 3.500 M K+; M NO − =
= 2.000 M NO3−
3
1.0000 L
1.0000 L
M CrO
=
2−
4
0.750 mol
= 0.750 M CrO42−; M Ag + = 0 M (the limiting reagent)
1.0000 L
CHAPTER 5
GASES
Questions
20.
Molecules in the condensed phases (liquids and solids) are very close together. Molecules in
the gaseous phase are very far apart. A sample of gas is mostly empty space. Therefore, one
would expect 1 mole of H2O(g) to occupy a huge volume as compared to 1 mole of H2O(l).
21.
The column of water would have to be 13.6 times taller than a column of mercury. When the
pressure of the column of liquid standing on the surface of the liquid is equal to the pressure
of air on the rest of the surface of the liquid, then the height of the column of liquid is a
measure of atmospheric pressure. Because water is 13.6 times less dense than mercury, the
column of water must be 13.6 times longer than that of mercury to match the force exerted by
the columns of liquid standing on the surface.
22.
A bag of potato chips is a constant-pressure container. The volume of the bag increases or
decreases in order to keep the internal pressure equal to the external (atmospheric) pressure.
The volume of the bag increased because the external pressure decreased. This seems
reasonable as atmospheric pressure is lower at higher altitudes than at sea level. We ignored n
(moles) as a possibility because the question said to concentrate on external conditions. It is
possible that a chemical reaction occurred that would increase the number of gas molecules
inside the bag. This would result in a larger volume for the bag of potato chips. The last
factor to consider is temperature. During ski season, one would expect the temperature of
Lake Tahoe to be colder than Los Angeles. A decrease in T would result in a decrease in the
volume of the potato chip bag. This is the exact opposite of what actually happened, so
apparently the temperature effect is not dominant.
23.
The P versus 1/V plot is incorrect. The plot should be linear with positive slope and a yintercept of zero. PV = k, so P = k(1/V). This is in the form of the straight-line equation y =
mx + b. The y-axis is pressure, and the x-axis is 1/V.
24.
The decrease in temperature causes the balloon to contract (V and T are directly related).
Because weather balloons do expand, the effect of the decrease in pressure must be dominant.
25.
d = (molar mass)P/RT; density is directly proportional to the molar mass of a gas. Helium,
with the smallest molar mass of all the noble gases, will have the smallest density.
26.
Rigid container: As temperature is increased, the gas molecules move with a faster average
velocity. This results in more frequent and more forceful collisions, resulting in an increase in
pressure. Density = mass/volume; the moles of gas are constant, and the volume of the
container is constant, so density in this case must be temperature-independent (density is
constant).
139
140
CHAPTER 5
GASES
Flexible container: The flexible container is a constant-pressure container. Therefore, the
final internal pressure will be unaffected by an increase in temperature. The density of the
gas, however, will be affected because the container volume is affected. As T increases, there
is an immediate increase in P inside the container. The container expands its volume to
reduce the internal pressure back to the external pressure. We have the same mass of gas in a
larger volume. Gas density will decrease in the flexible container as T increases.
27.
At STP (T = 273.2 K and P = 1.000 atm), the volume of 1.000 mol of gas is:
V=
nRT
=
P
1.000 mol ×
0.08206 L atm
× 273.2 K
K mol
= 22.42 L
1.000 atm
At STP, the volume of 1.000 mole of any gas is 22.42 L, assuming the gas behaves ideally.
Therefore, the molar volume of He(g) and N2(g) at STP both equal 22.42 L/mol. If the
temperature increases to 25.0°C (298.2 K), the volume of 1.000 mole of a gas will be larger
than 22.42 L/mole because molar volume is directly related to the temperature at constant
pressure. If 1.000 mole of a gas is collected over water at a total pressure of 1.000 atm, the
partial pressure of the collected gas will be less than 1.000 atm because water vapor is present
(Ptotal = Pgas + PH 2 O ). At some partial pressure below 1.000 atm, the volume of 1.000 mole of
a gas will be larger than 22.42 L/mol because molar volume is inversely related to the
pressure at constant temperature.
28.
For the first diagram, there is a total volume of 3X after the stopcock is open. The six total
gas particles will be equally distributed (on average) over the entire volume (3X). So per X
volume, there will be two gas particles. Your first drawing should have four gas particles in
the 2X volume flask and two gas particles in the X volume flask.
Applying Boyle’s law, the pressure in the two flasks after the stopcock is opened is:
P1V1 = P2V2, P2 =
P1V1
P × 2X
2
= 1
= P1
V2
3
3X
The final pressure in both flasks will be two-thirds that of the initial pressure in the left flask.
For the second diagram, there is a total volume of 2X after the stopcock is opened. The gas
particles will be equally distributed (on average) so that your drawing should have three gas
particles in each flask. The final pressure is:
P2 =
P1V1
P
P ×X
= 1
= 1
V2
2
2X
The final pressure in both flasks will be one-half that of the initial pressure in the left flask.
29.
No; at any nonzero Kelvin temperature, there is a distribution of kinetic energies. Similarly,
there is a distribution of velocities at any nonzero Kelvin temperature. The reason there is a
distribution of kinetic energies at any specific temperature is because there is a distribution of
velocities for any gas sample at any specific temperature.
CHAPTER 5
30.
GASES
141
a. Containers ii, iv, vi, and viii have volumes twice those of containers i, iii, v, and vii.
Containers iii, iv, vii, and viii have twice the number of molecules present than
containers i, ii, v, and vi. The container with the lowest pressure will be the one that has
the fewest moles of gas present in the largest volume (containers ii and vi both have the
lowest P). The smallest container with the most moles of gas present will have the
highest pressure (containers iii and vii both have the highest P). All the other containers
(i, iv, v, and viii) will have the same pressure between the two extremes. The order is ii =
vi < i = iv = v = viii < iii = vii.
b. All have the same average kinetic energy because the temperature is the same in each
container. Only temperature determines the average kinetic energy.
c. The least dense gas will be in container ii because it has the fewest of the lighter Ne
atoms present in the largest volume. Container vii has the most dense gas because the
largest number of the heavier Ar atoms are present in the smallest volume. To determine
the ordering for the other containers, we will calculate the relative density of each. In the
table below, m1 equals the mass of Ne in container i, V1 equals the volume of container i,
and d1 equals the density of the gas in container i.
Container
mass,
volume
density




 volume 
mass
i
m1, V1
m1
V1
= d1
ii
m1, 2V1
m1
1
= d1
2V1
2
iii
2m 1, V1
2 m1
V1
= 2d1
iv
2m 1, 2V1
2 m1
2 V1
= d1
v
2m 1, V1
2 m1
V1
= 2d1
vi
2m 1, 2V1
2 m1
2 V1
= d1
vii
4m 1, V1
4 m1
V1
= 4d1
viii
4m 1, 2V1
4 m1
2 V1
= 2d1
From the table, the order of gas density is ii < i = iv = vi < iii = v = viii < vii.
d. µrms = (3RT/M)1/2; the root mean square velocity only depends on the temperature and the
molar mass. Because T is constant, the heavier argon molecules will have a slower root
mean square velocity than the neon molecules. The order is v = vi = vii = viii < i = ii = iii
= iv.
31.
2 NH3(g) → N2(g) + 3 H2(g); as reactants are converted into products, we go from 2 moles
of gaseous reactants to 4 moles of gaseous products (1 mol N2 + 3 mol H2). Because the
moles of gas doubles as reactants are converted into products, the volume of the gases will
double (at constant P and T).
 RT 
PV = nRT, P = 
n = (constant)n; pressure is directly related to n at constant T and V.
 V 
As the reaction occurs, the moles of gas will double, so the pressure will double. Because 1
o
mole of N2 is produced for every 2 moles of NH3 reacted, PN 2 = (1/2) PNH
. Owing to the 3 :
3
o
2 mole ratio in the balanced equation, PH 2 = (3/2) PNH3 .
o
o
o
. As we said earlier, the total
Note: Ptotal = PH 2 + PN 2 = (3/2) PNH
+ (1/2) PNH
= 2PNH
3
3
3
pressure will double from the initial pressure of NH3 as reactants are completely converted
into products.
142
CHAPTER 5
GASES
32.
Statements a, c, and e are true. For statement b, if temperature is constant, then the average
kinetic energy will be constant no matter what the identity of the gas (KEave = 3/2 RT). For
statement d, as T increases, the average velocity of the gas particles increases. When gas
particles are moving faster, the effect of interparticle interactions is minimized. For statement
f, the KMT predicts that P is directly related to T at constant V and n. As T increases, the gas
molecules move faster, on average, resulting in more frequent and more forceful collisions.
This leads to an increase in P.
33.
The values of a are: H2,
0.244 atm L2
mol 2
; CO2, 3.59; N2, 1.39; CH4, 2.25
Because a is a measure of intermolecular attractions, the attractions are greatest for CO2.
34.
The van der Waals constant b is a measure of the size of the molecule. Thus C3H8 should
have the largest value of b because it has the largest molar mass (size).
35.
PV = nRT; Figure 5.6 is illustrating how well Boyle’s law works. Boyle’s law studies the
pressure-volume relationship for a gas at constant moles of gas (n) and constant temperature
(T). At constant n and T, the PV product for an ideal gas equals a constant value of nRT, no
matter what the pressure of the gas. Figure 5.6 plots the PV product versus P for three
different gases. The ideal value for the PV product is shown with a dotted line at about a
value of 22.41 L atm. From the plot, it looks like the plot for Ne is closest to the dotted line,
so we can conclude that of the three gases in the plot, Ne behaves most ideally. The O2 plot
is also fairly close to the dotted line, so O2 also behaves fairly ideally. CO2, on the other
hand, has a plot farthest from the ideal plot; hence CO2 behaves least ideally.
36.
Dalton’s law of partial pressures holds if the total pressure of a mixture of gases depends only
on the total moles of gas particles present and not on the identity of the gases in the mixtures.
If the total pressure of a mixture of gases were to depend on the identities of the gases, then
each gas would behave differently at a certain set of conditions, and determining the pressure
of a mixture of gases would be very difficult. All ideal gases are assumed volumeless and are
assumed to exert no forces among the individual gas particles. Only in this scenario can
Dalton’s law of partial pressure hold true for an ideal gas. If gas particles did have a volume
and/or did exert forces among themselves, then each gas, with its own identity and size,
would behave differently. This is not observed for ideal gases.
Exercises
Pressure
37.
a. 4.8 atm ×
760 mm Hg
= 3.6 × 103 mm Hg
atm
b. 3.6 × 103 mm Hg ×
c. 4.8 atm ×
1.013 × 105 Pa
= 4.9 × 105 Pa
atm
d. 4.8 atm ×
1 torr
mm Hg
= 3.6 × 103 torr
14.7 psi
= 71 psi
atm
CHAPTER 5
38.
39.
a.
GASES
2200 psi ×
143
1 atm
= 150 atm
14.7 psi
b. 150 atm ×
1.013 × 105 Pa
1 MPa
= 15 MPa
×
atm
1 × 10 6 Pa
c. 150 atm ×
760 torr
= 1.1 × 105 torr
atm
6.5 cm ×
10 mm
1 atm
= 65 mm Hg = 65 torr; 65 torr ×
= 8.6 × 10 −2 atm
cm
760 torr
8.6 × 10 −2 atm ×
1.013 × 105 Pa
= 8.7 × 103 Pa
atm
1 atm
2.54 cm 10 mm
= 508 mm Hg = 508 torr; 508 torr ×
= 0.668 atm
×
in
cm
760 torr
40.
20.0 in Hg ×
41.
If the levels of mercury in each arm of the manometer are equal, then the pressure in the flask
is equal to atmospheric pressure. When they are unequal, the difference in height in
millimeters will be equal to the difference in pressure in millimeters of mercury between the
flask and the atmosphere. Which level is higher will tell us whether the pressure in the flask
is less than or greater than atmospheric.
a. Pflask < Patm; Pflask = 760. − 118 = 642 torr
642 torr ×
1 atm
= 0.845 atm
760 torr
0.845 atm ×
1.013 × 105 Pa
= 8.56 × 104 Pa
atm
b. Pflask > Patm; Pflask = 760. torr + 215 torr = 975 torr
c.
42.
975 torr ×
1 atm
= 1.28 atm
760 torr
1.28 atm ×
1.013 × 105 Pa
= 1.30 × 105 Pa
atm
Pflask = 635 − 118 = 517 torr; Pflask = 635 + 215 = 850. torr
a. The pressure is proportional to the mass of the fluid. The mass is proportional to the
volume of the column of fluid (or to the height of the column assuming the area of the
column of fluid is constant).
144
CHAPTER 5
d = density =
V=
GASES
mass
; in this case, the volume of silicon oil will be the same as the
volume volume of mercury in Exercise 41.
m Hg
m d
m
m
; VHg = Voil;
= oil , m oil = Hg oil
d
d Hg
d oil
d Hg
Because P is proportional to the mass of liquid:
d
Poil = PHg  oil
 d Hg


 = PHg  1.30  = (0.0956)PHg

 13.6 

This conversion applies only to the column of silicon oil.
Pflask = 760. torr − (0.0956 × 118) torr = 760. − 11.3 = 749 torr
749 torr ×
1 atm
1.013 × 105 Pa
= 0.986 atm; 0.986 atm ×
= 9.99 × 104 Pa
760 torr
atm
Pflask = 760. torr + (0.0956 × 215) torr = 760. + 20.6 = 781 torr
781 torr ×
1 atm
1.013 × 105 Pa
= 1.03 atm; 1.03 atm ×
= 1.04 × 105 Pa
atm
760 torr
b. If we are measuring the same pressure, the height of the silicon oil column would be 13.6
÷ 1.30 = 10.5 times the height of a mercury column. The advantage of using a less dense
fluid than mercury is in measuring small pressures. The height difference measured will
be larger for the less dense fluid. Thus the measurement will be more precise.
Gas Laws
43.
At constant n and T, PV = nRT = constant, so P1V1 = P2V2; at sea level, P = 1.00 atm
= 760. mm Hg.
V2 =
P1 V1
760. mm × 2.0 L
=
= 3.0 L
500. mm Hg
P2
The balloon will burst at this pressure because the volume must expand beyond the 2.5 L
limit of the balloon.
Note: To solve this problem, we did not have to convert the pressure units into atm; the units
of mm Hg canceled each other. In general, only convert units if you have to. Whenever the
gas constant R is not used to solve a problem, pressure and volume units must only be
consistent and not necessarily in units of atm and L. The exception is temperature, which
must always be converted to the Kelvin scale.
CHAPTER 5
44.
GASES
145
The pressure exerted on the balloon is constant, and the moles of gas present is constant.
From Charles’s law, V1/T1 = V2/T2 at constant P and n.
V2 =
V1 T2
700. mL × 100. K
=
= 239 mL
(273.2 + 20.0) K
T1
As expected, as temperature decreases, the volume decreases.
45.
At constant T and P, Avogadro’s law holds (V ∝ n).
V1
V
Vn
20. L × 0.50 mol
= 0.89 mol
= 2 , n2 = 2 1 =
n1
n2
V1
11.2 L
As expected, as V increases, n increases.
46.
As NO2 is converted completely into N2O4, the moles of gas present will decrease by one-half
(from the 2 : 1 mole ratio in the balanced equation). Using Avogadro’s law:
V1
n
V
1
= 2 , V2 = V1 × 2 = 25.0 mL ×
= 12.5 mL
2
n1
n1
n2
N2O4(g) will occupy one-half the original volume of NO2(g). This is expected because the
moles of gas present decrease by one-half when NO2 is converted into N2O4.
47.
a. PV = nRT, V =
b. PV = nRT, n =
c. PV = nRT, T =
d. PV = nRT, P =
48.
nRT
=
P
0.08206 L atm
× (155 + 273) K
K mol
= 14.0 L
5.00 atm
PV
0.300 atm × 2.00 L
=
= 4.72 × 10 −2 mol
0.08206 L atm
RT
× 155 K
K mol
PV
4.47 atm × 25.0 L
=
0.08206 L atm
nR
2.01 mol ×
K mol
nRT
=
V
a. P = 7.74 × 103 Pa ×
PV = nRT, n =
2.00 mol ×
10.5 mol ×
1 atm
1.013 × 105 Pa
= 678 K = 405°C
0.08206 L atm
× (273 + 75) K
K mol
= 133 atm
2.25 L
= 0.0764 atm; T = 25 + 273 = 298 K
PV
0.0764 atm × 0.0122 L
=
= 3.81 × 10 −5 mol
0
.
08206
L
atm
RT
× 298 K
K mol
146
CHAPTER 5
b. PV = nRT, P =
nRT
c. V =
=
P
nRT
=
V
0.421 mol ×
0.08206 L atm
× 223 K
K mol
= 179 atm
0.0430 L
0.08206 L atm
× (331 + 273) K
K mol
= 3.6 L
1 atm
455 torr ×
760 torr
4.4 × 10 − 2 mol ×


1 atm
 745 mm Hg ×
 × 11.2 L
760 mm Hg 
PV

=
= 334 K = 61°C
d. T =
0.08206 L atm
nR
0.401 mol ×
K mol
49.
50.
n=
2.70 atm × 200.0 L
PV
=
= 22.2 mol
0.08206 L atm
RT
× (273 + 24) K
K mol
For He: 22.2 mol ×
4.003 g He
= 88.9 g He
mol
For H2: 22.2 mol ×
2.016 g H 2
= 44.8 g H2
mol
PV
1.00 atm × 6.0 L
=
= 0.25 mol air
0.08206 L atm
RT
× 298 K
K mol
a.
n =
b.
n =
1.97 atm × 6.0 L
= 0.48 mol air
0.08206 L atm/K mol × 298 K
c.
n =
0.296 atm × 6.0 L
= 0.11 mol air
0.08206 L atm/K mol × 200. K
Air is indeed “thinner” at high elevations.
51.
52.
PV = nRT, n =
14.5 atm × (75.0 × 10 −3 L)
PV
= 0.0449 mol O2
=
0.08206 L atm
RT
× 295 K
K mol

1 mol  0.08206 L atm
 0.60 g ×
×
× (273 + 22) K
32.00 g 
K mol
nRT

P =
=
= 0.091 atm
V
5.0 L
GASES
CHAPTER 5
53.
GASES
a. PV = nRT; 175 g Ar ×
T=
0.050 mL ×
nRT
=
V=
P
55.
1 mol Ar
= 4.38 mol Ar
39.95 g Ar
PV
10.0 atm × 2.50 L
=
= 69.6 K
0.08206 L atm
nR
4.38 mol ×
K mol
b. PV = nRT, P =
54.
147
nRT
=
V
0.08206 L atm
× 255 K
K mol
= 32.3 atm
2.50 L
1 mol O 2
1.149 g
= 1.8 × 10 −3 mol O2
×
mL
32.00 g
1.8 × 10 −3 mol ×
For a gas at two conditions:
Because V is constant:
n2 =
4.38 mol ×
0.08206 L atm
× 310. K
K mol
= 4.6 × 10 −2 L = 46 mL
1.0 atm
P1V1
PV
= 2 2
n1T1
n 2 T2
P1
P
nPT
= 2 , n2 = 1 2 1
n1T1
n 2 T2
P1T2
1.50 mol × 800. torr × 298 K
= 2.77 mol
400. torr × 323 K
Moles of gas added = n2 – n1 = 2.77 – 1.50 = 1.27 mol
For two-condition problems, units for P and V just need to be the same units for both conditions, not necessarily atm and L. The unit conversions from other P or V units would cancel
when applied to both conditions. However, temperature always must be converted to the
Kelvin scale. The temperature conversions between other units and Kelvin will not cancel
each other.
56.
PV = nRT, n is constant.
V2 = (1.040)V1,
P2 =
57.
PV
PV
PV
= nR = constant, 1 1 = 2 2
T
T1
T2
V1
1.000
=
1.040
V2
P1V1T2
1.000 (273 + 58) K
= 75 psi ×
= 82 psi
×
1.040 (273 + 19) K
V2 T1
P1V1 P2 V2
; all gases are assumed to follow the ideal gas law. The
=
n1T1 n 2 T2
identity of the gas in container B is unimportant as long as we know the moles of gas present.
At two conditions:
148
CHAPTER 5
GASES
PB
V n T
1.0 L × 2.0 mol × 560. K
= 2.0
= A B B =
PA
VB n A TA
2.0 L × 1.0 mol × 280. K
The pressure of the gas in container B is twice the pressure of the gas in container A.
58.
The pressure is doubled so P2 = 2P1 and the absolute temperature is halved so T2 = ½T1 (or T1
= 2T2). The moles of gas did not change, so n2 = n1. The volume effect of these changes is:
P1 V1
P V
V
Pn T
PT
P × T2
= 2 2, 2 = 1 2 2 = 1 2 = 1
= 1/4
n 1 T1
n 2 T2 V 1
P2 n 1 T1
P2 T1
2P1 × 2T2
The volume of the gas decreases by a factor of four when the pressure is doubled and the
absolute temperature is halved.
59.
a. At constant n and V,
b.
TP
P1 P2
6.50 atm
, T2 = 1 2 = 273 K ×
= 161 K
=
P1
11.0 atm
T1 T2
c. T2 =
60.
P1 P2
318 K
PT
, P2 = 1 2 = 11.0 atm ×
= 12.8 atm
=
T1 T2
T1
273 K
T1 P2
25.0 atm
= 273 K ×
= 620. K
P1
11.0 atm
Because the container is flexible, P is assumed constant. The moles of gas present are also
constant.
P1V1
PV V
V
= 2 2 , 1 = 2 ; Vsphere = 4/3 πr3
n1T1
n 2 T2 T1
T2
61.
V2 =
V1T2
4/3 π (1.00 cm) 3 × 361 K
, 4/3 π(r2 ) 3 =
T1
280. K
r23 =
361 K
= 1.29, r2 = (1.29)1/3 = 1.09 cm = radius of sphere after heating
280. K
PV
P V
PV
= nR = constant, 1 1 = 2 2
T1
T2
T
P2 =
62.
5.0 × 10 2 mL (273 + 820.) K
P1 V1 T2
×
= 710. torr ×
= 5.1 × 104 torr
25 mL
(273 + 30.) K
V2 T1
PV = nRT,
nT
V
n T
nT
= = constant, 1 1 = 2 2 ; moles × molar mass = mass
P1
P2
P
R
mass 2 × T2
n1 (molar mass)T1
n (molar mass)T2 mass1 × T1
,
= 2
=
P1
P2
P1
P2
CHAPTER 5
GASES
Mass2 =
63.
mass1 × T1P2
1.00 × 103 g × 291 K × 650. psi
=
= 309 g
T2 P1
299 K × 2050. psi
PV = nRT, n is constant.
V2 = 1.00 L ×
64.
149
PV
P V
PV
VPT
= nR = constant, 1 1 = 2 2 , V2 = 1 1 2
T
T1
T2
V2T1
760.torr
(273 − 31) K
= 2.82 L; ΔV = 2.82 - 1.00 = 1.82 L
×
220. torr (273 + 23) K
PV = nRT, P is constant.
nT P
nT
n T
= = constant, 1 1 = 2 2
V1
V2
V
R
294 K 4.20 × 103 m 3
n2
TV
= 0.921
×
= 1 2 =
n1
T2 V1
335 K 4.00 × 103 m 3
Gas Density, Molar Mass, and Reaction Stoichiometry
65.
STP: T = 273 K and P = 1.00 atm; at STP, the molar volume of a gas is 22.42 L.
2.00 L O2 ×
1 mol O 2
4 mol Al 26.98 g Al
= 3.21 g Al
×
×
mol Al
3 mol O 2
22.42 L
Note: We could also solve this problem using PV = nRT, where n O 2 = PV/RT. You don’t
have to memorize 22.42 L/mol at STP.
66.
CO2(s) → CO2(g); 4.00 g CO2 ×
1 mol CO 2
= 9.09 × 10 −2 mol CO2
44.01 g CO 2
At STP, the molar volume of a gas is 22.42 L. 9.09 × 10 −2 mol CO2 ×
67.
2 NaN3(s) → 2 Na(s) + 3 N2(g)
n N2 =
PV
1.00 atm × 70.0 L
=
= 3.12 mol N2 needed to fill air bag.
0
.
08206
L atm
RT
× 273 K
K mol
Mass NaN3 reacted = 3.12 mol N2 ×
68.
22.42 L
= 2.04 L
mol CO 2
2 mol NaN 3
65.02 g NaN 3
= 135 g NaN3
×
3 mol N 2
mol NaN 3
Because the solution is 50.0% H2O2 by mass, the mass of H2O2 decomposed is 125/2 =
62.5 g.
62.5 g H2O2 ×
1 mol O 2
1 mol H 2 O 2
= 0.919 mol O2
×
2 mol H 2 O 2
34.02 g H 2 O 2
150
CHAPTER 5
nRT
V=
=
P
69.
GASES
0.08206 L atm
× 300. K
K mol
= 23.0 L O2
1 atm
746 torr ×
760 torr
0.919 mol ×
3

1L 
 100 cm 
1.0 atm × 4800 m 3 × 

 ×
1000 cm 3 
 m 

PV

n H2 =
=
= 2.1 × 105 mol
0
.
08206
L
atm
RT
× 273 K
K mol
2.1 × 105 mol H2 is in the balloon. This is 80.% of the total amount of H2 that had to be
generated:
0.80(total mol H2) = 2.1 × 105, total mol H2 = 2.6 × 105 mol
70.
2.6 × 105 mol H2 ×
1 mol Fe 55.85 g Fe
= 1.5 × 107 g Fe
×
mol H 2
mol Fe
2.6 × 105 mol H2 ×
1 mol H 2SO 4
98.09 g H 2SO 4 100 g reagent
×
×
mol H 2
mol H 2SO 4
98 g H 2SO 4
= 2.6 × 107 g of 98% sulfuric acid
5.00 g S ×
1 mol S
= 0.156 mol S
32.07 g
0.156 mol S will react with 0.156 mol O2 to produce 0.156 mol SO2. More O2 is required to
convert SO2 into SO3.
0.156 mol SO2 ×
1 mol O 2
= 0.0780 mol O2
2 mol SO 2
Total mol O2 reacted = 0.156 + 0.0780 = 0.234 mol O2
nRT
V=
=
P
71.
0.234 mol ×
0.08206 L atm
× 623 K
K mol
= 2.28 L O2
5.25 atm
Kr(g) + 2 Cl2(g) → KrCl4(s); nKr =
PV
0.500 atm × 15.0 L
= 0.147 mol Kr
=
0.08206 L atm
RT
× 623 K
K mol
We could do the same calculation for Cl2. However, the only variable that changed is the
pressure. Because the partial pressure of Cl2 is triple that of Kr, moles of Cl2 = 3(0.147) =
0.441 mol Cl2. The balanced equation requires 2 moles of Cl2 to react with every mole of Kr.
However, we actually have three times as many moles of Cl2 as we have of Kr. So Cl2 is in
excess and Kr is the limiting reagent.
CHAPTER 5
GASES
0.147 mol Kr ×
72.
151
1 mol KrCl 4
225.60 g KrCl 4
= 33.2 g KrCl4
×
mol Kr
mol Kr
PV = nRT, V and T are constant.
P1
P
P
n
= 2, 2 = 2
n1
n 2 P1
n1
Let's calculate the partial pressure of C3H3N that can be produced from each of the starting
materials assuming each reactant is limiting. The reactant that produces the smallest amount
of product will run out first and is the limiting reagent.
PC 3 H 3 N = 0.500 MPa ×
2 MPa C 3 H 3 N
= 0.800 MPa if NH3 is limiting
2 MPa NH 3
PC3H 3 N = 0.800 MPa ×
PC3H 3 N = 1.500 MPa ×
2 MPa C3H3 N
= 0.500 MPa if C3H6 is limiting
2 MPa C3H 6
2 MPa C 3 H 3 N
= 1.000 MPa if O2 is limiting
3 MPa O 2
C3H6 is limiting. Although more product could be produced from NH3 and O2, there is only
enough C3H6 to produce 0.500 MPa of C3H3N. The partial pressure of C3H3N in atmospheres
after the reaction is:
0.500 × 106 Pa ×
n =
PV
4.94 atm × 150. L
= 30.3 mol C3H3N
=
0.08206 L atm
RT
× 298 K
K mol
30.3 mol ×
73.
1 atm
= 4.94 atm
1.013 × 105 Pa
53.06 g
= 1.61 × 103 g C3H3N can be produced.
mol
CH3OH + 3/2 O2 → CO2 + 2 H2O or 2 CH3OH(l) + 3 O2(g) → 2 CO2(g) + 4 H2O(g)
50.0 mL ×
n O2 =
0.850 g
1 mol
×
= 1.33 mol CH3OH(l) available
mL
32.04 g
PV
2.00 atm × 22.8 L
=
= 1.85 mol O2 available
0
.
08206
L atm
RT
× 300. K
K mol
Assuming CH3OH is limiting:
1.33 mol CH3OH ×
4 mol H 2 O
= 2.66 mol H2O
2 mol CH 3OH
152
CHAPTER 5
GASES
Assuming O2 is limiting:
1.85 mol O2 ×
4 mol H 2 O
= 2.47 mol H2O
3 mol O 2
Because the O2 reactant produces the smaller quantity of H2O, O2 is limiting and 2.47 mol of
H2O can be produced.
74.
For ammonia (in 1 minute):
n NH 3 =
PNH 3 × VNH 3
RT
=
90. atm × 500. L
= 1.1 × 103 mol NH3
0.08206 L atm
× 496 K
K mol
NH3 flows into the reactor at a rate of 1.1 × 103 mol/min.
For CO2 (in 1 minute):
n CO 2 =
PCO 2 × VCO 2
RT
=
45 atm × 600. L
= 6.6 × 102 mol CO2
0.08206 L atm
× 496 K
K mol
CO2 flows into the reactor at 6.6 × 102 mol/min.
If NH3 is limiting:
1.1 × 103 mol NH 3
1 mol urea
60.06 g urea
×
×
= 3.3 × 104 g urea/min
min
2 mol NH 3
mol urea
If CO2 is limiting:
660 mol CO 2 1 mol urea 60.06 g urea
×
×
= 4.0 × 104 g urea/min
min
mol CO 2
mol urea
Because the NH3 reactant produces the smaller quantity of product, NH3 is limiting and
3.3 × 104 g urea/min can be formed.
75.
a. CH4(g) + NH3(g) + O2(g) → HCN(g) + H2O(g); balancing H first, then O, gives:
CH4 + NH3 +
3
2
O 2 → HCN + 3 H2O or 2 CH4(g) + 2 NH3(g) + 3 O2(g) →
2 HCN(g) + 6 H2O(g)
b. PV = nRT, T and P constant;
V1
V
V
n
= 2, 1 = 1
n1
n 2 V2
n2
The volumes are all measured at constant T and P, so the volumes of gas present are
directly proportional to the moles of gas present (Avogadro’s law). Because Avogadro’s
law applies, the balanced reaction gives mole relationships as well as volume
relationships.
CHAPTER 5
GASES
153
If CH4 is limiting: 20.0 L CH4 ×
2 L HCN
= 20.0 L HCN
2 L CH 4
If NH3 is limiting: 20.0 L NH3 ×
2 L HCN
= 20.0 L HCN
2 L NH 3
If O2 is limiting: 20.0 L O2 ×
2 L HCN
= 13.3 L HCN
3 L O2
O2 produces the smallest quantity of product, so O2 is limiting and 13.3 L HCN can be
produced.
76.
From the balanced equation, ethene reacts with hydrogen in a 1 : 1 mole ratio. Because T and
P are constant, a greater volume of H2 and thus more moles of H2 are flowing into the
reaction container than moles of ethene. So ethene is the limiting reagent.
In 1 minute:
n C2H 4 =
PV
25.0 atm × 1000. L
=
= 532 mol C2H4 reacted
0
.
08206
L atm
RT
× 573 K
K mol
Theoretical yield =
Percent yield =
77.
Molar mass =
1 kg
532 mol C2 H 4 1 mol C2 H 6 30.07 g C2 H 6
×
×
×
1000 g
mol C2 H 6
min
mol C2 H 4
= 16.0 kg C2H6/min
15.0 kg/min
× 100 = 93.8%
16.0 kg/min
dRT
, where d = density of gas in units of g/L.
P
3.164 g/L ×
Molar mass =
0.08206 L atm
× 273.2 K
K mol
= 70.98 g/mol
1.000 atm
The gas is diatomic, so the average atomic mass = 70.93/2 = 35.47 u. From the periodic table,
this is chlorine, and the identity of the gas is Cl2.
78.
P × (molar mass) = dRT, d =
mass
mass
× RT
, P × (molar mass) =
volume
V
0.08206 L atm
0.800 g ×
× 373 K
mass × RT
K mol
Molar mass =
= 96.9 g/mol
=
1 atm
PV
(750. torr ×
) × 0.256 L
760 torr
154
CHAPTER 5
GASES
96.9
= 2.00; molecular formula is C2H2Cl2.
48.5

1 atm 
 745 torr ×
 × 352.0 g/mol
760 torr 
P × (molar mass)

=
=
= 12.6 g/L
0.08206 L atm
RT
× 333 K
K mol
Mass of CHCl ≈ 12.0 + 1.0 + 35.5 = 48.5 g/mol;
79.
d UF6
80.
d = P × (molar mass)/RT; we need to determine the average molar mass of air. We get this
by using the mole fraction information to determine the weighted value for the molar mass. If
we have 1.000 mol of air:
average molar mass = 0.78 mol N2 ×
28.02 g N 2
32.00 g O 2
+ 0.21 mol O2 ×
mol O 2
mol N 2
+ 0.010 mol Ar ×
dair =
1.00 atm × 29 g/mol
0.08206 L atm
× 273 K
K mol
39.95 g Ar
= 28.98 = 29 g
mol Ar
= 1.3 g/L
Partial Pressure
81.
The container has 5 He atoms, 3 Ne atoms, and 2 Ar atoms for a total of 10 atoms. The mole
fractions of the various gases will be equal to the molecule fractions.
χHe =
5 He atoms
3 Ne atoms
= 0.50; χNe =
= 0.30
10 total atoms
10 total atoms
χAr = 1.00 – 0.50 – 0.30 = 0.20
PHe = χHe × Ptotal = 0.50(1.00 atm) = 0.50 atm
PNe = χNe × PTotal = 0.30(1.00atm) = 0.30 atm
PAr = 1.00 atm – 0.50 atm – 0.30 atm = 0.20 atm
82.
a. There are 6 He atoms and 4 Ne atoms, and each flask has the same volume. The He flask
has 1.5 times as many atoms of gas present as the Ne flask, so the pressure in the He flask
will be 1.5 times greater (assuming a constant temperature).
b. Because the flask volumes are the same, your drawing should have the various atoms
equally distributed between the two flasks. So each flask should have 3 He atoms and 2
Ne atoms.
CHAPTER 5
GASES
155
c. After the stopcock is opened, each flask will have 5 total atoms and the pressures will be
equal. If six atoms of He gave an initial pressure of PHe, initial, then 5 total atoms will have
a pressure of 5/6 × PHe, initial.
Using similar reasoning, 4 atoms of Ne gave an initial pressure of PNe, initial, so 5 total
atoms will have a pressure of 5/4 × PNe, initial. Summarizing:
Pfinal =
5
5
PHe, initial = PNe, initial
6
4
d. For the partial pressures, treat each gas separately. For helium, when the stopcock is
opened, the six atoms of gas are now distributed over a larger volume. To solve for the
final partial pressures, use Boyle’s law for each gas.
For He: P2 =
PHe, initial
P1V1
X
= PHe, initial ×
=
V2
2X
2
The partial pressure of helium is exactly halved. The same result occurs with neon so
that when the volume is doubled, the partial pressure is halved. Summarizing:
PHe, final =
83.
PCO 2
PHe, initial
2
; PNe, final =
PNe, initial
2

1 mol  0.08206 L atm
 ×
 7.8 g ×
× 300. K
K mol
44.01 g 
nRT

=
=
= 1.1 atm
4.0 L
V
With air present, the partial pressure of CO2 will still be 1.1 atm. The total pressure will be
the sum of the partial pressures, Ptotal = PCO 2 + Pair.

1 atm 
 = 1.1 + 0.97 = 2.1 atm
Ptotal = 1.1 atm +  740 torr ×
760 torr 

84.
n H 2 = 1.00 g H2 ×
PH 2 =
PHe =
n H 2 × RT
V
1 mol H 2
1 mol He
= 0.496 mol H2; n He = 1.00 g He ×
4.003 g He
2.016 g H 2
= 0.250 mol He
0.496 mol ×
=
0.08206 L atm
× (273 + 27) K
K mol
= 12.2 atm
1.00 L
n He × RT
= 6.15 atm; Ptotal = PH 2 + PHe = 12.2 atm + 6.15 atm = 18.4 atm
V
156
85.
CHAPTER 5
GASES
Treat each gas separately and determine how the partial pressure of each gas changes when
the container volume increases. Once the partial pressures of H2 and N2 are determined, the
total pressure will be the sum of these two partial pressures. At constant n and T, the
relationship P1V1 = P2V2 holds for each gas.
For H2: P2 =
P1V1
2.00 L
= 475 torr ×
= 317 torr
V2
3.00 L
760 torr
1.00 L
= 0.0667 atm; 0.0667 atm ×
= 50.7 torr
atm
3.00 L
For N2: P2 = 0.200 atm ×
Ptotal = PH 2 + PN 2 = 317 + 50.7 = 368 torr
86.
For H2: P2 =
P1V1
2.00 L
= 360. torr ×
= 240. torr
V2
3.00 L
Ptotal = PH 2 + PN 2 , PN 2 = Ptotal − PH 2 = 320. torr − 240. torr = 80. torr
For N2: P1 =
87.
P2 V2
3.00 L
= 80. torr ×
= 240 torr
1.00 L
V1
P1V1 = P2V2; the total volume is 1.00 L + 1.00 L + 2.00 L = 4.00 L.
For He: P2 =
P1 V1
1.00 L
= 200. torr ×
= 50.0 torr He
V2
4.00 L
For Ne: P2 = 0.400 atm ×
For Ar: P2 = 24.0 kPa ×
760 torr
1.00 L
= 0.100 atm; 0.100 atm ×
= 76.0 torr Ne
4.00 L
atm
2.00 L
760 torr
1 atm
= 12.0 kPa; 12.0 kPa ×
×
4.00 L
atm
101.3 kPa
= 90.0 torr Ar
Ptotal = 50.0 + 76.0 + 90.0 = 216.0 torr
88.
We can use the ideal gas law to calculate the partial pressure of each gas or to calculate the
total pressure. There will be less math if we calculate the total pressure from the ideal gas
law.
n O 2 = 1.5 × 102 mg O2 ×
1 mol O 2
1g
= 4.7 × 10−3 mol O2
×
32.00 g O 2
1000 mg
n NH3 = 5.0 × 1021 molecules NH3 ×
1 mol NH 3
6.022 × 10
23
molecules NH 3
= 8.3 × 10−3 mol NH3
ntotal = n N 2 + n O 2 + n NH 3 = 5.0 × 10−2 + 4.7 × 10−3 + 8.3 × 10−3 = 6.3 × 10−2 mol total
CHAPTER 5
GASES
Ptotal =
157
n total × RT
=
V
6.3 × 10 − 2 mol ×
n N2
PN 2 = χ N 2 × Ptotal , χ N 2 =
PO 2 =
89.
4.7 × 10 −3
6.3 × 10 − 2
n total
; PN 2 =
0.08206 L atm
× 273 K
K mol
= 1.4 atm
1.0 L
5.0 × 10 −2 mol
6.3 × 10 − 2 mol
× 1.4 atm = 0.10 atm; PNH3 =
a. Mole fraction CH4 = χ CH 4 =
PCH 4
Ptotal
=
× 1.4 atm = 1.1 atm
8.3 × 10 −3
6.3 × 10 − 2
× 1.4 atm = 0.18 atm
0.175 atm
= 0.412
0.175 atm + 0.250 atm
χ O 2 = 1.000 − 0.412 = 0.588
b. PV = nRT, ntotal =
c.
n CH 4
χ CH 4 =
n total
Ptotal × V
RT
=
0.425 atm × 10.5 L
= 0.161 mol
0.08206 L atm
× 338 K
K mol
, n CH 4 = χ CH 4 × ntotal = 0.412 × 0.161 mol = 6.63 × 10 −2 mol CH4
6.63 × 10 −2 mol CH4 ×
16.04 g CH 4
= 1.06 g CH4
mol CH 4
n O 2 = 0.588 × 0.161 mol = 9.47 × 10 −2 mol O2; 9.47 × mol O2 ×
32.00 g O 2
mol O 2
= 3.03 g O2
90.
52.5 g O2 ×
χ O2 =
1 mol O 2
1 mol CO 2
= 1.64 mol O2; 65.1 g CO2 ×
= 1.48 mol CO2
32.00 g O 2
44.01 g CO 2
n O2
n total
=
1.64 mol
= 0.526
(1.64 + 1.48) mol
PO 2 = χ O 2 × Ptotal = 0.526 × 9.21 atm = 4.84 atm
PCO 2 = 9.21 – 4.84 = 4.37 atm
91.
Ptotal = PH 2 + PH 2O , 1.032 atm = PH 2 + 32 torr ×
n H2 =
PH 2 V
RT
=
1 atm
, PH 2 = 1.032 − 0.042 = 0.990 atm
760 torr
0.990 atm × 0.240 L
= 9.56 × 10 −3 mol H2
0.08206 L atm
× 303 K
K mol
158
CHAPTER 5
9.56 × 10 −3 mol H2 ×
92.
GASES
1 mol Zn 65.38 g Zn
= 0.625 g Zn
×
mol H 2
mol Zn
To calculate the volume of gas, we can use Ptotal and ntotal (V = ntotal RT/Ptotal), or we can use
PHe and nHe (V = nHeRT/PHe). Because n H 2O is unknown, we will use PHe and nHe.
PHe + PH 2 O = 1.00 atm = 760. torr, PHe + 23.8 torr = 760. torr, PHe = 736 torr
nHe = 0.586 g ×
n RT
V = He
=
PHe
93.
1 mol
= 0.146 mol He
4.003 g
0.08206 L atm
× 298 K
K mol
= 3.69 L
1 atm
736 torr ×
760 torr
0.146 mol ×
2 NaClO3(s) → 2 NaCl(s) + 3 O2(g)
Ptotal = PO 2 + PH 2 O , PO 2 = Ptotal − PH 2 O = 734 torr − 19.8 torr = 714 torr

1 atm 
 714 torr ×
 × 0.0572 L
PO 2 × V
760 torr 

n O2 =
=
= 2.22 × 10 −3 mol O2
0.08206 L atm
RT
× (273 + 22) K
K mol
Mass NaClO3 decomposed = 2.22 × 10 −3 mol O2 ×
Mass % NaClO3 =
94.
2 mol NaClO3 106.44 g NaClO3
×
3 mol O 2
mol NaClO3
= 0.158 g NaClO3
0.158 g
× 100 = 18.0%
0.8765 g
10.10 atm − 7.62 atm = 2.48 atm is the pressure of the amount of F2 reacted.
PV = nRT, V and T are constant.
P
P
P
P
n
= constant, 1 = 2 or 1 = 1
n1
n2
P2
n2
n
2.48 atm
Moles F2 reacted
=
= 2.00; so Xe + 2 F2 → XeF4
Moles Xe reacted
1.24 atm
95.
Because P and T are constant, V and n are directly proportional. The balanced equation
requires 2 L of H2 to react with 1 L of CO (2 : 1 volume ratio due to 2 : 1 mole ratio in the
balanced equation). If in 1 minute all 16.0 L of H2 react, only 8.0 L of CO are required to
react with it. Because we have 25.0 L of CO present in that 1 minute, CO is in excess and H2
is the limiting reactant. The volume of CH3OH produced at STP will be one-half the volume
of H2 reacted due to the 1 : 2 mole ratio in the balanced equation. In 1 minute, 16.0 L/2 =
8.00 L CH3OH is produced (theoretical yield).
CHAPTER 5
GASES
n CH 3OH =
159
PV
1.00 atm × 8.00 L
=
= 0.357 mol CH3OH in 1 minute
0.08206 L atm
RT
× 273 K
K mol
0.357 mol CH3OH ×
96.
32.04 g CH 3OH
= 11.4 g CH3OH (theoretical yield per minute)
mol CH 3OH
Percent yield =
actual yield
5.30 g
× 100 =
× 100 = 46.5% yield
theoretica l yield
11.4 g
750. mL juice ×
12 mL C2 H 5OH
= 90. mL C2H5OH present
100 mL juice
90. mL C2H5OH ×
1 mol C 2 H 5OH
0.79 g C 2 H 5OH
2 mol CO 2
×
×
= 1.5 mol CO2
2 mol C 2 H 5OH
46.07 g C 2 H 5OH
mL C 2 H 5OH
The CO2 will occupy (825 − 750. =) 75 mL not occupied by the liquid (headspace).
PCO 2 =
n CO 2 RT
V
1.5 mol ×
=
0.08206 L atm
× 298 K
K mol
= 490 atm
75 × 10 −3 L
Actually, enough CO2 will dissolve in the wine to lower the pressure of CO2 to a much more
reasonable value.
97.
2 HN3(g) → 3 N2(g) + H2(g); at constant V and T, P is directly proportional to n. In the
reaction, we go from 2 moles of gaseous reactants to 4 moles of gaseous products. Because
moles doubled, the final pressure will double (Ptotal = 6.0 atm). Similarly, from the 2 : 1 mole
ratio between HN3 and H2, the partial pressure of H2 will be 3.0/2 = 1.5 atm. The partial
pressure of N2 will be (3/2)3.0 atm = 4.5 atm. This is from the 2 : 3 mole ratio between HN3
and N2.
98.
2 SO2(g) + O2(g) → 2 SO3(g); because P and T are constant, volume ratios will equal mole
ratios (Vf/Vi = nf/ni). Let x = mol SO2 = mol O2 present initially. From the balanced
equation, 2 mol of SO2 react for every 1 mol of O2 that reacts. Because we have equal moles
of SO2 and O2 present initially, and because SO2 is used up twice as fast as O2, SO2 is the
limiting reagent. Therefore, no SO2 will be present after the reaction goes to completion.
However, excess O2(g) will be present as well as the SO3(g) produced.
Mol O2 reacted = x mol SO2 ×
1 mol O 2
= x/2 mol O2
2 mol SO 2
Mol O2 remaining = x mol O2 initially − x/2 mol O2 reacted = x/2 mol O2
Mol SO3 produced = x mol SO2 ×
2 mol SO 3
= x mol SO3
2 mol SO 2
160
CHAPTER 5
GASES
Total moles gas initially = x mol SO2 + x mol O2 = 2x
Total moles gas after reaction = x/2 mol O2 + x mol SO3 = (3/2)x = (1.5)x
nf
V
(1.5) x 1.5
= f =
=
= 0.75; Vf/Vi = 0.75 : l or 3 : 4
ni
Vi
2x
2
The volume of the reaction container shrinks to 75% of the initial volume.
99.
150 g (CH3)2N2H2 ×
PN 2 =
nRT
=
V
1 mol (CH 3 ) 2 N 2 H 2
3 mol N 2
= 7.5 mol N2 produced
×
60.10 g
mol (CH 3 ) 2 N 2 H 2
7.5 mol ×
0.08206 L atm
× 400. K
K mol
= 0.98 atm
250 L
We could do a similar calculation for PH 2O and PCO 2 and then calculate Ptotal (= PN 2 + PH 2O
+ PCO 2 ) . Or we can recognize that 9 total moles of gaseous products form for every mole of
(CH3)2N2H2 reacted (from the balanced equation given in the problem). This is three times the
moles of N2 produced. Therefore, Ptotal will be three times larger than PN 2 .
Ptotal = 3 × PN 2 = 3 × 0.98 atm = 2.9 atm.
100.
The partial pressure of CO2 that reacted is 740. − 390. = 350. torr. Thus the number of moles
of CO2 that react is given by:
350.
atm × 3.00 L
PV
760
= 5.75 × 10−2 mol CO2
n=
=
0.08206 L atm
RT
× 293 K
K mol
5.75 × 10−2 mol CO2 ×
Mass % MgO =
1 mol MgO 40.31 g MgO
= 2.32 g MgO
×
1 mol CO 2
mol MgO
2.32 g
× 100 = 81.4% MgO
2.85 g
Kinetic Molecular Theory and Real Gases
101.
KEavg = (3/2)RT; the average kinetic energy depends only on temperature. At each temperature, CH4 and N2 will have the same average KE. For energy units of joules (J), use R =
8.3145 J/K•mol. To determine average KE per molecule, divide the molar KEavg by Avogadro’s number, 6.022 × 1023 molecules/mol.
At 273 K: KEavg =
3 8.3145 J
× 273 K = 3.40 × 103 J/mol = 5.65 × 10−21 J/molecule
×
2
K mol
CHAPTER 5
GASES
161
3 8.3145 J
× 546 K = 6.81 × 103 J/mol = 1.13 × 10−20 J/molecule
×
2
K mol
At 546 K: KEavg =
102.
nAr =
n CH 4
n CH 4
228 g
= 5.71 mol Ar; χ CH 4 =
= 0.650 =
39.95 g/mol
n CH + n Ar
n CH + 5.71
4
4
0.650( n CH 4 + 5.71) = n CH 4 , 3.71 = (0.350)n CH 4 , n CH 4 = 10.6 mol CH4
KEavg =
3
2
RT for 1 mole of gas
KEtotal = (10.6 + 5.71) mol × 3/2 × 8.3145 J/K•mol × 298 K = 6.06 × 104 J = 60.6 kJ
1/ 2
103.
 3 RT 
µrms = 

 M 
, where R =
8.3145 J
and M = molar mass in kg.
K mol
For CH4, M = 1.604 × 10−2 kg, and for N2, M = 2.802 × 10−2 kg.
For CH4 at 273 K: µrms
1/ 2
8.3145 J

× 273 K
3 ×
K mol

=
 1.604 × 10 − 2 kg/mol








= 652 m/s
Similarly, µrms for CH4 at 546 K is 921 m/s.
For N2 at 273 K: µrms
8.3145 J

× 273 K
3 ×
K mol

=
 2.802 × 10 − 2 kg/mol


1/ 2






= 493 m/s
Similarly, for N2 at 546 K, µrms = 697 m/s.
1/ 2
104.
µrms
 3 RT 
=

 M 
;
μ UF6
μ He
 3 RTUF6

 M UF
6
= 
 3 RTHe

 M
He









1/ 2
1/ 2
 M He TUF6
=
 M UF THe
6





1/ 2
We want the root mean square velocities to be equal, and this occurs when:
M He TUF6 = M UF6 THe
The ratio of the temperatures is:
TUF6
THe
=
M UF6
M He
=
352.0
= 87.93
4.003
The heavier UF6 molecules would need a temperature 87.93 times that of the He atoms in
order for the root mean square velocities to be equal.
162
105.
106.
CHAPTER 5
GASES
The number of gas particles is constant, so at constant moles of gas, either a temperature
change or a pressure change results in the smaller volume. If the temperature is constant, an
increase in the external pressure would cause the volume to decrease. Gases are mostly
empty space so gases are easily compressible.
If the pressure is constant, a decrease in temperature would cause the volume to decrease. As
the temperature is lowered, the gas particles move with a slower average velocity and don’t
collide with the container walls as frequently and as forcefully. As a result, the internal
pressure decreases. In order to keep the pressure constant, the volume of the container must
decrease in order to increase the gas particle collisions per unit area.
In this situation, the volume has increased by a factor of two. One way to double the volume
of a container at constant pressure and temperature is to double the number of moles of gas
particles present. As gas particles are added, more collisions per unit area occur and the
internal pressure increases. In order to keep the pressure constant, the container volume must
increase.
Another way to double the volume of a container at constant pressure and moles of gas is to
double the absolute temperature. As temperature increases, the gas molecules collide more
frequently with the walls of the container. In order to keep pressure constant, the container
volume must increase.
The last variable which can be changed is pressure. If the external pressure exerted on the
container is halved, the volume will double (assuming constant temperature and moles). As
the external pressure applied is reduced, the volume of the container must increase in order to
equalize the higher internal pressure with the lower external applied pressure.
107.
a
b
c
d
Avg. KE
increase
decrease
same (KE ∝ T)
Avg. velocity
increase
decrease
same (
1
mv2 = KE ∝ T)
same
same
2
Wall coll. freq
increase
decrease
increase
increase
Average kinetic energy and average velocity depend on T. As T increases, both average
kinetic energy and average velocity increase. At constant T, both average kinetic energy and
average velocity are constant. The collision frequency is proportional to the average velocity
(as velocity increases, it takes less time to move to the next collision) and to the quantity n/V
(as molecules per volume increase, collision frequency increases).
108.
V, T, and P are all constant, so n must be constant. Because we have equal moles of gas in
each container, gas B molecules must be heavier than gas A molecules.
a. Both gas samples have the same number of molecules present (n is constant).
b. Because T is constant, KEavg must be the same for both gases [KEavg = (3/2)RT].
c. The lighter gas A molecules will have the faster average velocity.
CHAPTER 5
GASES
163
d. The heavier gas B molecules do collide more forcefully, but gas A molecules, with the
faster average velocity, collide more frequently. The end result is that P is constant
between the two containers.
109.
a. They will all have the same average kinetic energy because they are all at the same temperature [KEavg = (3/2)RT].
b. Flask C; H2 has the smallest molar mass. At constant T, the lighter molecules have the
faster average velocity. This must be true for the average kinetic energies to be the same.
110.
a. All the gases have the same average kinetic energy since they are all at the same
temperature [KEavg = (3/2)RT].
b. At constant T, the lighter the gas molecule, the faster the average velocity [µavg ∝ µrms ∝
(1/M)1/2].
Xe (131.3 g/mol) < Cl2 (70.90 g/mol) < O2 (32.00 g/mol) < H2 (2.016 g/mol)
slowest
fastest
c. At constant T, the lighter H2 molecules have a faster average velocity than the heavier O2
molecules. As temperature increases, the average velocity of the gas molecules
increases. Separate samples of H2 and O2 can only have the same average velocities if
the temperature of the O2 sample is greater than the temperature of the H2 sample.
1/ 2
111.
Graham’s law of effusion:
M 
Rate1
=  2 
Rate 2
 M1 
Let Freon-12 = gas 1 and Freon-11 = gas 2:
1/ 2
1.07  137.4 

=
1.00  M1 
, 1.14 =
137.4
, M1 = 121 g/mol
M1
The molar mass of CF2Cl2 is equal to 121 g/mol, so Freon-12 is CF2Cl2.
112.
M 
Rate 1
=  2 
Rate 2
 M1 
24.0  16.04 

=
47.8  M 1 
1/ 2
M 
Rate1
=  2 
Rate 2
 M1 
24.0 mL
16.04 g
47.8 mL
; M1 = ?
; rate2 =
; M2 =
min
min
mol
1/ 2
1/ 2
113.
; rate1 =
= 0.502, 16.04 = (0.502)2 × M1, M1 =
rate (12 C17 O)
The relative rates of effusion of 12C16O to
1/ 2
1/ 2
 30.0 
=
,

12 18
rate ( C O)  29.0 
12
16.04
63.7 g
=
0.252
mol
Rate (12 C16 O)  30.0 
=

Rate (12 C18 O)  28.0 
= 1.02;
C17O to
12
C18O are 1.04 : 1.02 : 1.00.
= 1.04
164
CHAPTER 5
GASES
Advantage: CO2 isn't as toxic as CO.
Major disadvantages of using CO2 instead of CO:
1. Can get a mixture of oxygen isotopes in CO2.
2. Some species, for example, 12C16O18O and 12C17O2, would effuse (gaseously diffuse)
at about the same rate because the masses are about equal. Thus some species cannot
be separated from each other.
1/ 2
114.
Rate1  M 2 

=
Rate 2  M1 
, where M = molar mass; let gas (1) = He and gas (2) = Cl2.
Effusion rates in this problem are equal to the volume of gas that effuses per unit time
(L/min). Let t = time in the following expression.
1.0 L
1/ 2
4.5 min
t
 70.90 
= 
= 4.209, t = 19 min
 ,
1.0 L
4
.
003
4
.
5
min


t
115.
a. PV = nRT
nRT
=
P=
V
b.
0.5000 mol ×
0.08206 L atm
× (25.0 + 273.2) K
K mol
= 12.24 atm
1.0000 L
2

n 
2
2
P + a   (V − nb) = nRT; for N2: a = 1.39 atm L /mol and b = 0.0391 L/mol
 V  

2


 0.5000 
P
+
1
.
39


 atm  (1.0000 L − 0.5000 × 0.0391 L) = 12.24 L atm
 1.0000 


(P + 0.348 atm)(0.9805 L) = 12.24 L atm
P=
12.24 L atm
− 0.348 atm = 12.48 − 0.348 = 12.13 atm
0.9805 L
c. The ideal gas law is high by 0.11 atm, or
116.
0.11
× 100 = 0.91%.
12.13
a. PV = nRT
P=
nRT
=
V
0.5000 mol ×
0.08206 L atm
× 298.2 K
K mol
= 1.224 atm
10.000 L
CHAPTER 5
b.
GASES
165
2

n 
2
2
P + a   (V – nb) = nRT; for N2: a = 1.39 atm L /mol and b = 0.0391 L/mol
V
  

2


 0.5000 
P + 1.39
 atm  (10.000 L − 0.5000 × 0.0391 L) = 12.24 L atm
 10.000 


(P + 0.00348 atm)(10.000 L − 0.0196 L) = 12.24 L atm
P + 0.00348 atm =
12.24 L atm
= 1.226 atm, P = 1.226 − 0.00348 = 1.223 atm
9.980 L
c. The results agree to ±0.001 atm (0.08%).
d. In Exercise 115, the pressure is relatively high, and there is significant disagreement. In
Exercise 116, the pressure is around 1 atm, and both gas laws show better agreement.
The ideal gas law is valid at relatively low pressures.
Atmospheric Chemistry
117.
χHe = 5.24 × 10 −6 from Table 5.4. PHe = χHe × Ptotal = 5.24 × 10 −6 × 1.0 atm = 5.2 × 10 −6 atm
5.2 × 10 −6 atm
n
P
=
= 2.1 × 10 −7 mol He/L
=
0
.
08206
L
atm
V
RT
× 298 K
K mol
2.1 × 10 −7 mol
1L
6.022 × 10 23 atoms
= 1.3 × 1014 atoms He/cm3
×
×
L
mol
1000 cm 3
118.
At 15 km, T ≈ −60°C and P = 0.1 atm. Use
V2 =
119.
P1 V1
P V
= 2 2 since n is constant.
T1
T2
V1 P1 T2
1.0 L × 1.00 atm × 213 K
=7L
=
P2 T1
0.1 atm × 298 K
S(s) + O2(g) → SO2(g), combustion of coal
2 SO2(g) + O2(g) → 2 SO3(g), reaction with atmospheric O2
SO3(g) + H2O(l) → H2SO4(aq), reaction with atmospheric H2O
120.
H2SO4(aq) + CaCO3(s) → CaSO4(aq) + H2O(l) + CO2(g)
121.
a. If we have 1.0 × 106 L of air, then there are 3.0 × 102 L of CO.
166
CHAPTER 5
VCO
3.0 × 10 2
because V ∝ n; PCO =
× 628 torr = 0.19 torr
Vtotal
1.0 × 10 6
PCO = χCOPtotal; χCO =
b. nCO =
PCO V
;
RT
nCO
GASES
assuming 1.0 m3 air, 1 m3 = 1000 L:
0.19
atm × (1.0 × 103 L)
760
=
= 1.1 × 10−2 mol CO
0.08206 L atm
× 273 K
K mol
1.1 × 10−2 mol ×
6.02 × 10 23 molecules
= 6.6 × 1021 CO molecules in 1.0 m3 of air
mol
3
c.
122.
6.6 × 10 21 molecules  1 m 
6.6 × 1015 molecules CO


=
×
 100 cm 
m3
cm 3


For benzene:
89.6 × 10-9 g ×
Vbenzene =
1 mol
= 1.15 × 10−9 mol benzene
78.11 g
n benzene RT
=
P
Mixing ratio =
Or ppbv =
0.08206 L atm
× 296 K
K mol
= 2.84 × 10−8 L
1 atm
748 torr ×
760 torr
1.15 × 10 −9 mol ×
2.84 × 10 −8 L
× 106 = 9.47 × 10−3 ppmv
3.00 L
vol. of X × 109
2.84 × 10 −8 L
× 109 = 9.47 ppbv
=
total vol.
3.00 L
1.15 × 10 −9 mol benzene
1L
6.022 × 10 23 molecules
×
×
3.00 L
mol
1000 cm 3
= 2.31 × 1011 molecules benzene/cm3
For toluene:
153 × 10−9 g C7H8 ×
1 mol
= 1.66 × 10−9 mol toluene
92.13 g
n
RT
Vtoluene = toluene
=
P
0.08206 L atm
× 296 K
K mol
= 4.10 × 10−8 L
1 atm
748 torr ×
760 torr
1.66 × 10 −9 mol ×
CHAPTER 5
GASES
Mixing ratio =
167
4.10 × 10 −8 L
× 106 = 1.37 × 10−2 ppmv (or 13.7 ppbv)
3.00 L
1.66 × 10 −9 mol toluene
1L
6.022 × 10 23 molecules
×
×
3.00 L
mol
1000 cm 3
= 3.33 × 1011 molecules toluene/cm3
Additional Exercises
123.
a. PV = nRT
PV = constant
b. PV = nRT
c. PV = nRT
 nR 
P= 
 × T = const × T
 V 
P
PV
T
V
d. PV = nRT
e. P =
nR
constant
=
V
V
P = constant ×
P
V
V
T
PV = constant
P
 P 
T= 
 × V = const × V
 nR 
f.
PV = nRT
PV
= nR = constant
T
1
V
PV
T
1/V
P
Note: The equation for a straight line is y = mx + b, where y is the y-axis and x is the x-axis.
Any equation that has this form will produce a straight line with slope equal to m and a y
intercept equal to b. Plots b, c, and e have this straight-line form.
124.
At constant T and P, Avogadro’s law applies; that is, equal volumes contain equal moles of
molecules. In terms of balanced equations, we can say that mole ratios and volume ratios
between the various reactants and products will be equal to each other. Br2 + 3 F2 → 2 X; 2
moles of X must contain 2 moles of Br and 6 moles of F; X must have the formula BrF3 for a
balanced equation.
125.
14.1 × 102 in Hg•in3 ×
3
2.54 cm 10 mm
1 atm
1L
 2.54 cm 
×
×
× 
 ×
in
1 cm
760 mm  in 
1000 cm 3
168
CHAPTER 5
GASES
= 0.772 atm•L
Boyle’s law: PV = k, where k = nRT; from Example 5.3, the k values are around 22 atm•L.
Because k = nRT, we can assume that Boyle’s data and the Example 5.3 data were taken at
different temperatures and/or had different sample sizes (different moles).
126.
Mn(s) + x HCl(g) → MnClx(s) +
n H2 =
x
H2(g)
2
PV
0.951 atm × 3.22 L
=
= 0.100 mol H2
RT 0.08206 L atm × 373 K
K mol
Mol Cl in compound = mol HCl = 0.100 mol H2 ×
x mol Cl
= 0.200 mol Cl
x
mol H 2
2
0.200 mol Cl
0.200 mol Cl
Mol Cl
=
= 4.00
=
Mol Mn 2.747 g Mn × 1 mol Mn
0.05000 mol Mn
54.94 g Mn
The formula of compound is MnCl4.
127.
Assume some mass of the mixture. If we had 100.0 g of the gas, we would have 50.0 g He
and 50.0 g Xe.
χ He =
n He
n He + n Xe
50.0 g
12.5 mol He
4.003 g/mol
=
= 0.970
=
50.0 g
50.0 g
12
.
5
mol
He
0
.
381
mol
Xe
+
+
4.003 g/mol 131.3 g/mol
No matter what the initial mass of mixture is assumed, the mole fraction of helium will
always be 0.970.
PHe = χHePtotal = 0.970 × 600. torr = 582 torr; PXe = 600. − 582 = 18 torr
128.
Assuming 100.0 g of cyclopropane:
85.7 g C ×
1 mol C
= 7.14 mol C
12.01 g
14.3 g H ×
1 mol H
14.2
= 14.2 mol H;
= 1.99
1.008 g
7.14
The empirical formula for cyclopropane is CH2, which has an empirical mass ≈ 12.0 + 2(1.0)
= 14.0 g/mol.
CHAPTER 5
GASES
169
P × (molar mass) = dRT, molar mass =
dRT
=
P
1.88 g / L ×
0.08206 L atm
× 273 K
K mol
1.00 atm
= 42.1 g/mol
Because 42.1/14.0 ≈ 3.0, the molecular formula for cyclopropane is (CH2)× 3 = C3H6.
129.
Ptotal = PN 2 + PH 2O , PN 2 = 726 torr – 23.8 torr = 702 torr ×
n N2 =
PN 2 × V
RT
=
0.924 atm × 31.8 × 10 −3 L
= 1.20 × 10−3 mol N2
0.08206 L atm
× 298 K
K mol
Mass of N in compound = 1.20 × 10−3 mol N2 ×
Mass % N =
130.
1 atm
= 0.924 atm
760 torr
28.02 g N 2
= 3.36 × 10−2 g nitrogen
mol
3.36 × 10 −2 g
× 100 = 13.3% N
0.253 g
33.5 mg CO2 ×
12.01 mg C
9.14 mg
= 9.14 mg C; % C =
× 100 = 26.1% C
44.01 mg CO 2
35.0 mg
41.1 mg H2O ×
2.016 mg H
4.60 mg
= 4.60 mg H; % H =
× 100 = 13.1% H
35.0 mg
18.02 mg H 2 O
n N2
740.
atm × 35.6 × 10 −3 L
760
= 1.42 × 10−3 mol N2
=
=
0
.
08206
L
atm
RT
× 298 K
K mol
PN 2 V
1.42 × 10-3 mol N2 ×
Mass % N =
28.02 g N 2
= 3.98 × 10−2 g nitrogen = 39.8 mg nitrogen
mol N 2
39.8 mg
× 100 = 61.0% N
65.2 mg
Or we can get % N by difference: % N = 100.0 − (26.1 + 13.1) = 60.8%
Out of 100.0 g:
26.1 g C ×
1 mol
2.17
= 2.17 mol C;
= 1.00
12.01 g
2.17
13.1 g H ×
1 mol
13.0
= 13.0 mol H;
= 5.99
1.008 g
2.17
170
CHAPTER 5
60.8 g N ×
1 mol
4.34
= 4.34 mol N;
= 2.00; empirical formula is CH6N2.
14.01 g
2.17
1/ 2
Rate1
 M 
=

Rate 2
 39.95 
131.
GASES
=
26.4
= 1.07, M = (1.07)2 × 39.95 = 45.7 g/mol
24.6
Empirical formula mass of CH6N2 ≈ 12 + 6 + 28 = 46 g/mol. Thus the molecular formula is
also CH6N2.
We will apply Boyle’s law to solve. PV = nRT = constant, P1V1 = P2V2
Let condition (1) correspond to He from the tank that can be used to fill balloons. We must
leave 1.0 atm of He in the tank, so P1 = 200. − 1.00 = 199 atm and V1 = 15.0 L. Condition (2)
will correspond to the filled balloons with P2 = 1.00 atm and V2 = N(2.00 L), where N is the
number of filled balloons, each at a volume of 2.00 L.
199 atm × 15.0 L = 1.00 atm × N(2.00 L), N = 1492.5; we can't fill 0.5 of a balloon, so N =
1492 balloons or, to 3 significant figures, 1490 balloons.
132.
Mol of He removed =
PV 1.00 atm × 1.75 × 10 −3 L
=
= 7.16 × 10 −5 mol
0.08206 L atm
RT
× 298 K
K mol
In the original flask, 7.16 × 10 −5 mol of He exerted a partial pressure of 1.960 − 1.710
= 0.250 atm.
V=
133.
7.16 × 10 −5 mol × 0.08206 × 298 K
nRT
=
= 7.00 × 10 −3 L = 7.00 mL
0.250 atm
P
For O2, n and T are constant, so P1V1 = P2V2.
P1 =
P2 V2
1.94 L
= 785 torr ×
= 761 torr = PO 2
V1
2.00 L
Ptotal = PO 2 + PH 2 O , PH 2 O = 785 − 761 = 24 torr
134.
PV = nRT, V and T are constant.
P1
P
P
n
= 2 or 1 = 1
n1
n2
P2
n2
When V and T are constant, then pressure is directly proportional to moles of gas present, and
pressure ratios are identical to mole ratios.
At 25°C: 2 H2(g) + O2(g) → 2 H2O(l), H2O(l) is produced at 25°C.
The balanced equation requires 2 mol H2 for every mol O2 reacted. The same ratio (2 : 1)
holds true for pressure units. So if all 2.00 atm of H2 react, only 1.00 atm of O2 will react
with it. Because we have 3.00 atm of O2 present, oxygen is in excess and hydrogen is the
limiting reactant. The only gas present at 25°C after the reaction goes to completion will be
the excess O2.
CHAPTER 5
GASES
171
PO 2 (reacted) = 2.00 atm H2 ×
1 atm O 2
= 1.00 atm O2
2 atm H 2
PO 2 (excess) = PO 2 (initial) − PO 2 (reacted) = 3.00 atm - 1.00 atm = 2.00 atm O2 = Ptotal
At 125°C: 2 H2(g) + O2(g) → 2 H2O(g), H2O(g) is produced at 125°C.
The major difference in the problem at 125°C versus 25°C is that gaseous water is now a
product (instead of liquid H2O), which will increase the total pressure because an additional
gas is present. Note: For this problem, it is assumed that 2.00 atm of H2 and 3.00 atm of O2
are reacted at 125°C instead of 25°C.
PH 2O (produced) = 2.00 atm H2 ×
2 atm H 2 O
= 2.00 atm H2O
2 atm H 2
Ptotal = PO 2 (excess) + PH 2O (produced) = 2.00 atm O2 + 2.00 atm H2O = 4.00 atm = Ptotal
135.
1.00 × 103 kg Mo ×
1 mol Mo
1000 g
= 1.04 × 104 mol Mo
×
95.94 g Mo
kg
1.04 × 104 mol Mo ×
VO 2 =
n O 2 RT
P
1 mol MoO 3
7/2 mol O 2
= 3.64 × 104 mol O2
×
mol Mo
mol MoO 3
3.64 × 10 4 mol ×
=
8.66 × 105 L O2 ×
100 L air
= 4.1 × 106 L air
21 L O 2
1.04 × 104 mol Mo ×
3 mol H 2
= 3.12 × 104 mol H2
mol Mo
3.12 × 10 4 mol ×
VH 2 =
136.
For NH3: P2 =
For O2: P2 =
0.08206 L atm
× 290. K
K mol
= 8.66 × 105 L of O2
1.00 atm
0.08206 L atm
× 290. K
K mol
= 7.42 × 105 L of H2
1.00 atm
2.00 L
P1V1
= 0.500 atm ×
= 0.333 atm
V2
3.00 L
1.00 L
P1V1
= 1.50 atm ×
= 0.500 atm
V2
3.00 L
After the stopcock is opened, V and T will be constant, so P ∝ n.
Assuming NH3 is limiting: 0.333 atm NH3 ×
4 atm NO
= 0.333 atm NO
4 atm NH 3
172
CHAPTER 5
Assuming O2 is limiting: 0.500 atm O2 ×
GASES
4 atm NO
= 0.400 atm NO
5 atm O 2
NH3 produces the smaller amount of product, so NH3 is limiting and 0.333 atm of NO can be
produced.
137.
Out of 100.00 g of compound there are:
1 mol C
4.872
= 4.872 mol C;
= 2.001
12.01 g C
2.435
58.51 g C ×
7.37 g H ×
1 mol H
7.31
= 7.31 mol H;
= 3.00
1.008 g H
2.435
1 mol N
2.435
= 2.435 mol N;
= 1.000
14.01 g N
2.435
34.12 g N ×
The empirical formula is C2H3N.
1/ 2
M 
Rate1
=  2 
Rate 2
 M1 
1/ 2
 M2 
; let gas (1) = He; 3.20 = 

 4.003 
, M 2 = 41.0 g/mol
The empirical formula mass of C2H3N ≈ 2(12.0) + 3(1.0) + 1(14.0) = 41.0 g/mol. So the
molecular formula is also C2H3N.
138.
If Be3+, the formula is Be(C5H7O2)3 and molar mass ≈ 13.5 + 15(12) + 21(1) + 6(16)
= 311 g/mol. If Be2+, the formula is Be(C5H7O2)2 and molar mass ≈ 9.0 + 10(12) + 14(1) +
4(16) = 207 g/mol.
Data set I (molar mass = dRT/P and d = mass/V):
0.08206 L atm
× 286 K
mass × RT
K mol
= 209 g/mol
molar mass =
=
1 atm
PV
(765.2 torr ×
) × (22.6 × 10 −3 L)
760 torr
0.2022 g ×
Data set II:
0.08206 L atm
× 290. K
mass × RT
K mol
= 202 g/mol
molar mass =
=
1 atm
PV
(764.6 torr ×
) × (26.0 × 10 −3 L)
760 torr
0.2224 g ×
These results are close to the expected value of 207 g/mol for Be(C5H7O2)2. Thus we
conclude from these data that beryllium is a divalent element with an atomic weight (mass) of
9.0 u.
139.
0.2766 g CO2 ×
12.01 g C
7.548 × 10 −2 g
= 7.548 × 10 −2 g C; % C =
× 100 = 73.78% C
0.1023 g
44.01 g CO 2
CHAPTER 5
GASES
173
2.016 g H
1.11 × 10 −2 g
= 1.11 × 10 −2 g H; % H =
× 100 = 10.9% H
0.1023 g
18.02 g H 2 O
0.0991 g H2O ×
PV = nRT, n N 2 =
PV 1.00 atm × 27.6 × 10 −3 L
=
= 1.23 × 10 −3 mol N2
0
.
08206
L
atm
RT
× 273 K
K mol
1.23 × 10 −3 mol N2 ×
Mass % N =
28.02 g N 2
= 3.45 × 10 −2 g nitrogen
mol N 2
3.45 × 10 −2 g
× 100 = 7.14% N
0.4831 g
Mass % O = 100.00 − (73.78 + 10.9 + 7.14) = 8.2% O
Out of 100.00 g of compound, there are:
73.78 g C ×
10.9 g H ×
1 mol
1 mol
= 6.143 mol C; 7.14 g N ×
= 0.510 mol N
12.01 g
14.01 g
1 mol
1 mol
= 10.8 mol H; 8.2 g O ×
= 0.51 mol O
1.008 g
16.00 g
Dividing all values by 0.51 gives an empirical formula of C12H21NO.
4.02 g 0.08206 L atm
×
× 400. K
dRT
L
K mol
=
= 392 g/mol
Molar mass =
1 atm
P
256 torr ×
760 torr
Empirical formula mass of C12H21NO ≈ 195 g/mol;
392
≈2
195
Thus the molecular formula is C24H42N2O2.
140.
At constant T, the lighter the gas molecules, the faster the average velocity. Therefore, the
pressure will increase initially because the lighter H2 molecules will effuse into container A
faster than air will escape. However, the pressures will eventually equalize once the gases
have had time to mix thoroughly.
ChemWork Problems
The answers to the problems 141-148 (or a variation to these problems) are found in OWL. These
problems are also assignable in OWL.
Challenge Problems
149.
BaO(s) + CO2(g) → BaCO3(s); CaO(s) + CO2(g) → CaCO3(s)
174
CHAPTER 5
GASES
750.
atm × 1.50 L
Pi V
760
ni =
= initial moles of CO2 =
= 0.0595 mol CO2
0.08206 L atm
RT
× 303.2 K
K mol
230.
atm × 1.50 L
Pf V
760
nf =
= final moles of CO2 =
= 0.0182 mol CO2
0.08206 L atm
RT
× 303.2 K
K mol
0.0595 − 0.0182 = 0.0413 mol CO2 reacted
Because each metal reacts 1 : 1 with CO2, the mixture contains a total of 0.0413 mol of BaO
and CaO. The molar masses of BaO and CaO are 153.3 and 56.08 g/mol, respectively.
Let x = mass of BaO and y = mass of CaO, so:
x + y = 5.14 g and
x
y
= 0.0413 mol or x + (2.734)y = 6.33
+
153.3 56.08
Solving by simultaneous equations:
x + (2.734)y = 6.33
−x
−y = −5.14
(1.734)y = 1.19, y – 1.19/1.734 = 0.686
y = 0.686 g CaO and 5.14 − y = x = 4.45 g BaO
Mass % BaO =
150.
4.45 g BaO
× 100 = 86.6% BaO; %CaO = 100.0 − 86.6 = 13.4% CaO
5.14 g
Cr(s) + 3 HCl(aq) → CrCl3(aq) + 3/2 H2(g); Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g)

1 atm 
 750. torr ×
 × 0.225 L
760 torr 
PV

Mol H2 produced = n =
=
= 9.02 × 10−3 mol H2
0
.
08206
L
atm
RT
× (273 + 27) K
K mol
9.02 × 10−3 mol H2 = mol H2 from Cr reaction + mol H2 from Zn reaction
From the balanced equation: 9.02 × 10−3 mol H2 = mol Cr × (3/2) + mol Zn × 1
Let x = mass of Cr and y = mass of Zn, then:
x + y = 0.362 g and 9.02 × 10−3 =
(1.5) x
y
+
52.00
65.38
We have two equations and two unknowns. Solving by simultaneous equations:
9.02 × 10−3 = (0.02885)x + (0.01530)y
−0.01530 × 0.362 = −(0.01530)x − (0.01530)y
CHAPTER 5
GASES
175
3.48 × 10−3 = (0.01355)x,
x = mass of Cr =
y = mass of Zn = 0.362 g − 0.257 g = 0.105 g Zn; mass % Zn =
151.
3.48 × 10 −3
= 0.257 g
0.01355
0.105 g
× 100
0.362 g
= 29.0% Zn
Assuming 1.000 L of the hydrocarbon (CxHy), then the volume of products will be 4.000 L,
and the mass of products (H2O + CO2) will be:
1.391 g/L × 4.000 L = 5.564 g products
Mol CxHy = n C x H y =
Mol products = np =
PV
0.959 atm × 1.000 L
=
= 0.0392 mol
0
.
08206
L atm
RT
× 298 K
K mol
PV
1.51 atm × 4.000 L
=
= 0.196 mol
0
.
08206
L atm
RT
× 375 K
K mol
CxHy + oxygen → x CO2 + y/2 H2O
Setting up two equations:
(0.0392)x + 0.0392(y/2) = 0.196
(moles of products)
(0.0392)x(44.01 g/mol) + 0.0392(y/2)(18.02 g/mol) = 5.564 g
(mass of products)
Solving: x = 2 and y = 6, so the formula of the hydrocarbon is C2H6.
152.
a.
Let x = moles SO2 = moles O2 and z = moles He.
P • MM
, where MM = molar mass
RT
1.924 g/L =
1.000 atm × MM
, MMmixture = 43.13 g/mol
0.08206 L atm
× 273.2 K
K mol
Assuming 1.000 total moles of mixture is present, then: x + x + z = 1.000 and:
64.07 g/mol × x + 32.00 g/mol × x + 4.003 g/mol × z = 43.13 g
2x + z = 1.000 and (96.07)x + (4.003)z = 43.13
Solving: x = 0.4443 mol and z = 0.1114 mol
Thus: χHe = 0.1114 mol/1.000 mol = 0.1114
b.
2 SO2(g) + O2(g) → 2 SO3(g)
176
CHAPTER 5
GASES
Initially, assume 0.4443 mol SO2, 0.4443 mol O2, and 0.1114 mol He. Because SO2 is
limiting, we end up with 0.2222 mol O2, 0.4443 mol SO3, and 0.1114 mol He in the
gaseous product mixture. This gives ninitial = 1.0000 mol and nfinal = 0.7779 mol.
In a reaction, mass is constant. d =
mass
1
and V ∝ n at constant P and T, so d ∝ .
V
n
n initial
d
1.0000
 1.0000 
=
= final , d final = 
 × 1.924 g/L, dfinal = 2.473 g/L
n final
0.7779
d initial
 0.7779 
153.
a. The reaction is CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g).
PV = nRT,
PCH 4 VCH 4
PV
P V
= RT = constant,
= air air
n
n CH 4
n air
The balanced equation requires 2 mol O2 for every mole of CH4 that reacts. For three
times as much oxygen, we would need 6 mol O2 per mole of CH4 reacted (n O 2 = 6n CH 4 ).
Air is 21% mole percent O2, so n O 2 = (0.21)nair. Therefore, the moles of air we would
need to deliver the excess O2 are:
n air
n O 2 = (0.21)nair = 6n CH 4 , nair = 29n CH 4 ,
= 29
n CH 4
In 1 minute:
Vair = VCH 4 ×
PCH 4
n air
1.50 atm
= 200. L × 29 ×
= 8.7 × 103 L air/min
×
1.00 atm
n CH 4
Pair
b. If x mol of CH4 were reacted, then 6x mol O2 were added, producing (0.950)x mol CO2
and (0.050)x mol of CO. In addition, 2x mol H2O must be produced to balance the
hydrogens.
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g); CH4(g) + 3/2 O2(g) → CO(g) + 2 H2O(g)
Amount O2 reacted:
(0.950)x mol CO2 ×
2 mol O 2
= (1.90)x mol O2
mol CO 2
(0.050)x mol CO ×
1.5 mol O 2
= (0.075)x mol O2
mol CO
Amount of O2 left in reaction mixture = (6.00)x − (1.90)x − (0.075)x = (4.03)x mol O2
Amount of N2 = (6.00)x mol O2 ×
The reaction mixture contains:
79 mol N 2
= (22.6)x ≈ 23x mol N2
21 mol O 2
CHAPTER 5
GASES
177
(0.950)x mol CO2 + (0.050)x mol CO + (4.03)x mol O2 + (2.00)x mol H2O
+ 23x mol N2 = (30.)x mol of gas total
χ CO =
χ H 2O =
154.
(0.050) x
(0.950) x
(4.03) x
= 0.0017; χ CO 2 =
= 0.032; χ O 2 =
= 0.13
(30.) x
(30.) x
(30.) x
(2.00) x
= 0.067;
(30.) x
χ N2 =
23 x
= 0.77
(30.) x
The reactions are:
C(s) + 1/2 O2(g) → CO(g) and C(s) + O2(g) → CO2(g)
 RT 
PV = nRT, P = n 
 = n(constant)
 V 
Because the pressure has increased by 17.0%, the number of moles of gas has also increased
by 17.0%.
nfinal = (1.170)ninitial = 1.170(5.00) = 5.85 mol gas = n O 2 + n CO + n CO 2
n CO + n CO 2 = 5.00 (balancing moles of C). Solving by simultaneous equations:
n O 2 + n CO + n CO 2 = 5.85
− (n CO + n CO 2 = 5.00)
_______________________
n O2
= 0.85
If all C were converted to CO2, no O2 would be left. If all C were converted to CO, we would
get 5 mol CO and 2.5 mol excess O2 in the reaction mixture. In the final mixture, moles of
CO equals twice the moles of O2 present ( n CO = 2n O 2 ).
n CO = 2n O 2 = 1.70 mol CO; 1.70 + n CO 2 = 5.00, n CO 2 = 3.30 mol CO2
χ CO =
155.
3.30
1.70
= 0.564;
= 0.291; χ CO 2 =
5.85
5.85
a. Volume of hot air: V =
χ O2 =
4 3
4
πr = π(2.50 m) 3 = 65.4 m3
3
3
(Note: Radius = diameter/2 = 5.00/2 = 2.50 m)
3
0.85
= 0.145 ≈ 0.15
5.85
1L
 10 dm 
65.4 m3 × 
= 6.54 × 104 L
 ×
3
dm
 m 
178
CHAPTER 5
GASES

1 atm 
 745 torr ×
 × 6.54 × 10 4 L
760
torr
PV

n=
= 2.31 × 103 mol air
= 
0.08206 L atm
RT
× (273 + 65) K
K mol
Mass of hot air = 2.31 × 103 mol ×
29.0 g
= 6.70 × 104 g
mol
745
atm × 6.54 × 10 4 L
PV
760
= 2.66 × 103 mol air
Air displaced: n =
=
0.08206 L atm
RT
× (273 + 21) K
K mol
Mass of air displaced = 2.66 × 103 mol ×
29.0 g
= 7.71 × 104 g
mol
Lift = 7.71 × 104 g − 6.70 × 104 g = 1.01 × 104 g
b. Mass of air displaced is the same, 7.71 × 104 g. Moles of He in balloon will be the same
as moles of air displaced, 2.66 × 103 mol, because P, V, and T are the same.
Mass of He = 2.66 × 103 mol ×
4.003 g
= 1.06 × 104 g
mol
Lift = 7.71 × 104 g − 1.06 × 104 g = 6.65 × 104 g
630.
atm × (6.54 × 10 4 L)
PV
c. Hot air: n =
= 1.95 × 103 mol air
= 760
0.08206 L atm
RT
× 338 K
K mol
1.95 × 103 mol ×
29.0 g
= 5.66 × 104 g of hot air
mol
630.
atm × (6.54 × 10 4 L)
PV
= 2.25 × 103 mol air
Air displaced: n =
= 760
0.08206 L atm
RT
× 294 K
K mol
2.25 × 103 mol ×
29.0 g
= 6.53 × 104 g of air displaced
mol
Lift = 6.53 × 104 g − 5.66 × 104 g = 8.7 × 103 g
156.
a. When the balloon is heated, the balloon will expand (P and n remain constant). The mass
of the balloon is the same, but the volume increases, so the density of the argon in the
balloon decreases. When the density is less than that of air, the balloon will rise.
b. Assuming the balloon has no mass, when the density of the argon equals the density of
air, the balloon will float in air. Above this temperature, the balloon will rise.
CHAPTER 5
GASES
dair =
179
P • MM air
, where MMair = average molar mass of air
RT
MMair = 0.790 × 28.02 g/mol + 0.210 × 32.00 g/mol = 28.9 g/mol
dair =
1.00 atm × 28.9 g/mol
= 1.18 g/L
0.08206 L atm
× 298 K
K mol
dargon =
1.00 atm × 39.95 g/mol
= 1.18 g/L, T = 413 K
0.08206 L atm
×T
K mol
Heat the Ar above 413 K or 140.°C, and the balloon would float.
157.
a.
Average molar mass of air = 0.790 × 28.02 g/mol + 0.210 × 32.00 g/mol = 28.9 g/mol
Molar mass of helium = 4.003 g/mol
A given volume of air at a given set of conditions has a larger density than helium at
those conditions due to the larger average molar mass of air. We need to heat the air to a
temperature greater than 25°C in order to lower the air density (by driving air molecules
out of the hot air balloon) until the density is the same as that for helium (at 25°C and
1.00 atm).
b. To provide the same lift as the helium balloon (assume V = 1.00 L), the mass of air in the
hot air balloon (V = 1.00 L) must be the same as that in the helium balloon. Let MM =
molar mass:
P•MM = dRT, mass =
Mass air = 0.164 g =
MM • PV
; solving: mass He = 0.164 g
RT
28.9 g/mol × 1.00 atm × 1.00 L
0.08206 L atm
×T
K mol
T = 2150 K (a very high temperature)
158.

an 2 
an 2 V
an 3 b
P +
 × (V − nb) = nRT, PV +
= nRT
−
nbP
−

V 2 
V2
V2

PV +
an 2
an 3 b
= nRT
− nbP −
V
V2
At low P and high T, the molar volume of a gas will be relatively large. Thus the an2/V and
an3b/V2 terms become negligible at low P and high T because V is large. Because nb is the
actual volume of the gas molecules themselves, nb << V and the −nbP term will be negligible as compared to PV. Thus PV = nRT.
180
159.
CHAPTER 5
GASES
d = molar mass(P/RT); at constant P and T, the density of gas is directly proportional to the
molar mass of the gas. Thus the molar mass of the gas has a value which is 1.38 times that of
the molar mass of O2.
Molar mass = 1.38(32.00 g/mol) = 44.2 g/mol
Because H2O is produced when the unknown binary compound is combusted, the unknown
must contain hydrogen. Let AxHy be the formula for unknown compound.
Mol AxHy = 10.0 g AxHy ×
Mol H = 16.3 g H2O ×
1 mol A x H y
44.2 g
= 0.226 mol AxHy
1 mol H 2 O
2 mol H
= 1.81 mol H
×
18.02 g
mol H 2 O
1.81 mol H
= 8 mol H/mol AxHy; AxHy = AxH8
0.226 mol A x H y
The mass of the x moles of A in the AxH8 formula is:
44.2 g − 8(1.008 g) = 36.1 g
From the periodic table and by trial and error, some possibilities for AxH8 are ClH8, F2H8,
C3H8, and Be4H8. C3H8 and Be4H8 fit the data best, and because C3H8 (propane) is a known
substance, C3H8 is the best possible identity from the data in this problem.
160.
a. Initially PN 2 = PH 2 = 1.00 atm, and the total pressure is 2.00 atm (P total = PN 2 + PH 2 ). The
total pressure after reaction will also be 2.00 atm because we have a constant-pressure
container. Because V and T are constant before the reaction takes place, there must be
equal moles of N2 and H2 present initially. Let x = mol N2 = mol H2 that are present
initially. From the balanced equation, N2(g) + 3 H2(g) → 2 NH3(g), H2 will be limiting
because three times as many moles of H2 are required to react as compared to moles of
N2. After the reaction occurs, none of the H2 remains (it is the limiting reagent).
Mol NH3 produced = x mol H2 ×
Mol N2 reacted = x mol H2 ×
2 mol NH 3
= 2x/3
3 mol H 2
1 mol N 2
= x/3
3 mol H 2
Mol N2 remaining = x mol N2 present initially − x/3 mol N2 reacted = 2x/3 mol N2
After the reaction goes to completion, equal moles of N2(g) and NH3(g) are present
(2x/3). Because equal moles are present, the partial pressure of each gas must be equal
(PN 2 = PNH 3 ).
CHAPTER 5
GASES
181
Ptotal = 2.00 atm = PN 2 + PNH 3 ; solving: PN 2 = 1.00 atm = PNH 3
b. V ∝ n because P and T are constant. The moles of gas present initially are:
n N 2 + n H 2 = x + x = 2x mol
After reaction, the moles of gas present are:
2x 2x
n N 2 + n NH 3 =
= 4x/3 mol
+
3
3
4 x/ 3
2
Vafter
n
=
= after =
2x
3
Vinitial
n initial
The volume of the container will be two-thirds the original volume, so:
V = 2/3(15.0 L) = 10.0 L
Integrative Problems
161.
The redox equation must be balanced. Each uranium atom changes oxidation sates from +4
in UO2+ to +6 in UO22+ (a loss of two electrons for each uranium atom). Each nitrogen atom
changes oxidation states from +5 in NO3− to +2 in NO (a gain of three electrons for each
nitrogen atom). To balance the electrons transferred, we need two N atoms for every three U
atoms. The balanced equation is:
2 H+(aq) + 2 NO3−(aq) + 3 UO2+(aq) → 3 UO22+(aq) + 2 NO(g) + H2O(l)
nNO =
PV
1.5 atm × 0.255 L
= 0.015 mol NO
=
0
.
08206
L atm
RT
× 302 K
K mol
0.015 mol NO ×
162.
a. 156 mL ×
nHCl =
3 mol UO 2+
= 0.023 mol UO2+
2 mol NO
1.34 g
= 209 g HSiCl3 = actual yield of HSiCl3
mL
PV
10.0 atm × 15.0 L
= 5.93 mol HCl
=
0
.
08206
L atm
RT
× 308 K
K mol
5.93 mol HCl ×
1 mol HSiCl 3 135.45 g HSiCl 3
= 268 g HSiCl3
×
3 mol HCl
1 mol HSiCl 3
Percent yield =
actual yield
209 g
× 100 =
× 100 = 78.0%
theoretical yield
268 g
182
CHAPTER 5
b. 209 g HiSCl3 ×
GASES
1 mol HSiCl 3
1 mol SiH 4
= 0.386 mol SiH4
×
135.45 g HSiCl 3
4 mol HSiCl 3
This is the theoretical yield. If the percent yield is 93.1%, then the actual yield is:
0.386 mol SiH4 × 0.931 = 0.359 mol SiH4
VSiH 4 =
163.
nRT
=
P
0.359 mol ×
0.08206 L atm
× 308 K
K mol
10.0 atm
= 0.907 L = 907 mL SiH4
ThF4, 232.0 + 4(19.00) = 308.0 g/mL
d=
molar mass × P
308.0 g/mol × 2.5 atm
= 4.8 g/L
=
0.08206 L atm
RT
× (1680 + 273) K
K mol
The gas with the smaller molar mass will effuse faster. Molar mass of ThF4 = 308.0 g/mol;
molar mass of UF3 = 238.0 + 3(19.00) = 295.0 g/mol. Therefore, UF3 will effuse faster.
Rate of effusion of UF3
=
Rate of effusion of ThF4
molar mass of ThF4
=
molar mass of UF3
308.0 g/mol
= 1.02
295.0 g/mol
UF3 effuses 1.02 times faster than ThF4.
164.
The partial pressures can be determined by using the mole fractions.
Pmethane = Ptotal × χmethane = 1.44 atm × 0.915 = 1.32 atm; Pethane = 1.44 – 1.32 = 0.12 atm
Determining the number of moles of natural gas combusted:
nnatural gas =
PV
1.44 atm × 15.00 L
= 0.898 mol natural gas
=
0.08206 L atm
RT
× 293 K
K mol
nmethane = nnatural gas × χmethane = 0.898 mol × 0.915 = 0.822 mol methane
nethane = 0.898 − 0.822 = 0.076 mol ethane
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l);
0.822 mol CH4 ×
0.076 mol C2H6 ×
2 C2H6 + 7 O2(g) → 4 CO2(g) + 6 H2O(l)
2 mol H 2 O 18.02 g H 2 O
= 29.6 g H2O
×
1 mol CH 4
mol H 2 O
6 mol H 2 O 18.02 g H 2 O
= 4.1 g H2O
×
mol H 2 O
2 mol C 2 H 6
CHAPTER 5
GASES
183
The total mass of H2O produced = 29.6 g + 4.1 g = 33.7 g H2O.
Marathon Problem
165.
a. The formula of the compound AxBy depends on which gas is limiting, A2 or B2. We need
to determine both possible products. The procedure we will use is to assume one reactant
is limiting, and then determine what happens to the initial total moles of gas as it is
converted into the product. Because P and T are constant, volume ∝ n. Because mass is
conserved in a chemical reaction, any change in density must be due to a change in
volume of the container as the reaction goes to completion.
Density = d ∝
d
n
1
and V ∝ n, so: after = initial
d initial
n after
V
Assume the molecular formula of the product is AxBy where x and y are whole numbers.
First, let’s consider when A2 is limiting with x moles each of A2 and B2 in our equimolar
mixture. Note that the coefficient in front of AxBy in the equation must be 2 for a
balanced reaction.
x A2(g)
Initial
Change
Final
+ y B2(g)
x mol
−x mol
0
→
2 AxBy(g)
x mol
−y mol
(x − y) mol
0 mol
+2 mol
2 mol
d after
n
2x
= 1.50 = initial =
x− y +2
d initial
n after
(1.50)x − (1.50)y + 3.00 = 2x, 3.00 − (1.50)y = (0.50)x
Because x and y are whole numbers, y must be 1 because the above equation does not
allow y to be 2 or greater. When y = 1, x = 3 giving a formula of A3B if A2 is limiting.
Assuming B2 is limiting with y moles in the equimolar mixture:
Initial
Change
After
x A2(g)
+ y B2(g)
y
−x
y−x
y
−y
0
→
density after
n
2y
= 1.50 = initial =
density before
n after
y−x+2
2 AxBy(g)
0
+2
2
184
CHAPTER 5
GASES
Solving gives x = 1 and y = 3 for a molecular formula of AB3 when B2 is limiting.
b. In both possible products, the equations dictated that only one mole of either A or B had
to be present in the formula. Any number larger than 1 would not fit the data given in the
problem. Thus the two formulas determined are both molecular formulas and not just
empirical formulas.
CHAPTER 6
THERMOCHEMISTRY
Questions
11.
Path-dependent functions for a trip from Chicago to Denver are those quantities that depend
on the route taken. One can fly directly from Chicago to Denver, or one could fly from
Chicago to Atlanta to Los Angeles and then to Denver. Some path-dependent quantities are
miles traveled, fuel consumption of the airplane, time traveling, airplane snacks eaten, etc.
State functions are path-independent; they only depend on the initial and final states. Some
state functions for an airplane trip from Chicago to Denver would be longitude change,
latitude change, elevation change, and overall time zone change.
12.
Products have a lower potential energy than reactants when the bonds in the products are
stronger (on average) than in the reactants. This occurs generally in exothermic processes.
Products have a higher potential energy than reactants when the reactants have the stronger
bonds (on average). This is typified by endothermic reactions.
13.
2 C8H18(l) + 25 O2(g) → 16 CO2(g) + 18 H2O(g); the combustion of gasoline is exothermic
(as is typical of combustion reactions). For exothermic reactions, heat is released into the
surroundings giving a negative q value. To determine the sign of w, concentrate on the moles
of gaseous reactants versus the moles of gaseous products. In this combustion reaction, we go
from 25 moles of reactant gas molecules to 16 + 18 = 34 moles of product gas molecules. As
reactants are converted to products, an expansion will occur because the moles of gas
increase. When a gas expands, the system does work on the surroundings, and w is a negative
value.
14.
∆H = ∆E + P∆V at constant P; from the definition of enthalpy, the difference between ∆H
and ∆E, at constant P, is the quantity P∆V. Thus, when a system at constant P can do
pressure-volume work, then ∆H ≠ ∆E. When the system cannot do PV work, then ∆H = ∆E at
constant pressure. An important way to differentiate ∆H from ∆E is to concentrate on q, the
heat flow; the heat flow by a system at constant pressure equals ∆H, and the heat flow by a
system at constant volume equals ∆E.
15.
a. The ∆H value for a reaction is specific to the coefficients in the balanced equation. Because the coefficient in front of H2O is a 2, 891 kJ of heat is released when 2 mol of H2O
is produced. For 1 mol of H2O formed, 891/2 = 446 kJ of heat is released.
b. 891/2 = 446 kJ of heat released for each mol of O2 reacted.
16.
Use the coefficients in the balanced rection to determine the heat required for the various
quantities.
a. 1 mol Hg ×
90.7 kJ
= 90.7 kJ required
mol Hg
184
CHAPTER
6
THERMOCHEMISTRY
b. 1 mol O2 ×
185
90.7 kJ
= 181.4 kJ required
1 / 2 mol O 2
c. When an equation is reversed, ∆Hnew = −∆Hold. When an equation is multiplied by some
integer n, then ∆Hnew = n(∆Hold).
Hg(l) + 1/2 O2(g) → HgO(s)
2Hg(l) + O2(g) → 2HgO(s)
17.
∆H = −90.7 kJ
∆H = 2(−90.7 kJ) = −181.4 kJ; 181.4 kJ released
Given:
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l)
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
∆H = −891 kJ
∆H = −803 kJ
Using Hess’s law:
H2O(l) + 1/2 CO2(g) → 1/2 CH4(g) + O2(g)
1/2 CH4(g) + O2(g) → 1/2 CO2(g) + H2O(g)
H2O(l) → H2O(g)
∆H1 = −1/2(−891 kJ)
∆H2 = 1/2(−803 kJ)
∆H = ∆H1 + ∆H2 = 44 kJ
The enthalpy of vaporization of water is 44 kJ/mol.
Note: When an equation is reversed, the sign on ΔH is reversed. When the coefficients in a
balanced equation are multiplied by an integer, then the value of ΔH is multiplied by the
same integer.
18.
A state function is a function whose change depends only on the initial and final states and
not on how one got from the initial to the final state. An extensive property depends on the
amount of substance. Enthalpy changes for a reaction are path-independent, but they do
depend on the quantity of reactants consumed in the reaction. Therefore, enthalpy changes
are a state function and an extensive property.
19.
The zero point for ΔH of values are elements in their standard state. All substances are measured in relationship to this zero point.
20.
a. CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l)
∆H = ?
Utilizing Hess’s law:
Reactants → Standard State Elements ∆H = ∆Ha + ∆Hb = 75 + 0 = 75 kJ
Standard State Elements → Products
Reactants → Products
∆H = ∆Hc + ∆Hd = –394 – 572 = –966 kJ
∆H = 75 – 966 = –891 kJ
b. The standard enthalpy of formation for an element in its standard state is given a value of
zero. To assign standard enthalpy of formation values for all other substances, there
needs to be a reference point from which all enthalpy changes are determined. This
reference point is the elements in their standard state which is defined as the zero point.
186
CHAPTER 6
THERMOCHEMISTRY
So when using standard enthalpy values, a reaction is broken up into two steps. The first
step is to calculate the enthalpy change necessary to convert the reactants to the elements
in their standard state. The second step is to determine the enthalpy change that occurs
when the elements in their standard state go to form the products. When these two steps
are added together, the reference point (the elements in their standard state) cancels out
and we are left with the enthalpy change for the reaction.
c. This overall reaction is just the reverse of all the steps in the part a answer. So ∆H° =
+966 – 75 = 891 kJ. Products are first converted to the elements in their standard state
which requires 966 kJ of heat. Next, the elements in the standard states go to form the
original reactants [CH4(g) + 2 O2(g)] which has an enthalpy change of −75 kJ. All of the
signs are reversed because the entire process is reversed.
21.
No matter how insulated your thermos bottle, some heat will always escape into the
surroundings. If the temperature of the thermos bottle (the surroundings) is high, less heat
initially will escape from the coffee (the system); this results in your coffee staying hotter for
a longer period of time.
22.
From the photosynthesis reaction, CO2(g) is used by plants to convert water into glucose and
oxygen. If the plant population is significantly reduced, not as much CO2 will be consumed
in the photosynthesis reaction. As the CO2 levels of the atmosphere increase, the greenhouse
effect due to excess CO2 in the atmosphere will become worse.
23.
Fossil fuels contain carbon; the incomplete combustion of fossil fuels produces CO(g) instead
of CO2(g). This occurs when the amount of oxygen reacting is not sufficient to convert all the
carbon to CO2. Carbon monoxide is a poisonous gas to humans.
24.
Advantages: H2 burns cleanly (less pollution) and gives a lot of energy per gram of fuel.
Water as a source of hydrogen is abundant and cheap.
Disadvantages: Expensive and gas storage and safety issues
Exercises
Potential and Kinetic Energy
25.
KE =
1 2
1 kg m 2
mv ; convert mass and velocity to SI units. 1 J =
2
s2
Mass = 5.25 oz ×
Velocity =
1 lb
1 kg
= 0.149 kg
×
16 oz 2.205 lb
1.0 × 10 2 mi
1h
1 min 1760 yd
1m
45 m
×
×
×
×
=
h
60 min
60 s
mi
1.094 yd
s
1
1
 45 m 
KE =
mv2 =
× 0.149 kg × 

2
2
 s 
2
= 150 J
CHAPTER
6
26.
1
1
 1.0 m 
mv2 =
× 2.0 kg × 

2
2
 s 
KE =
THERMOCHEMISTRY
187
2
2
= 1.0 J; KE =
1
1
 2.0 m 
mv2 =
× 1.0 kg × 

2
2
 s 
= 2.0 J
The 1.0-kg object with a velocity of 2.0 m/s has the greater kinetic energy.
27.
a. Potential energy is energy due to position. Initially, ball A has a higher potential energy
than ball B because the position of ball A is higher than the position of ball B. In the
final position, ball B has the higher position so ball B has the higher potential energy.
b. As ball A rolled down the hill, some of the potential energy lost by A has been converted
to random motion of the components of the hill (frictional heating). The remainder of the
lost potential energy was added to B to initially increase its kinetic energy and then to
increase its potential energy.
28.
Ball A: PE = mgz = 2.00 kg ×
9.81 m
196 kg m 2
× 10.0 m =
= 196 J
2
s
s2
At point I: All this energy is transferred to ball B. All of B's energy is kinetic energy at this
point. Etotal = KE = 196 J. At point II, the sum of the total energy will equal
196 J.
At point II: PE = mgz = 4.00 kg ×
9.81 m
× 3.00 m = 118 J
s2
KE = Etotal − PE = 196 J − 118 J = 78 J
Heat and Work
29.
ΔE = q + w = 45 kJ + (−29 kJ) = 16 kJ
30.
ΔE = q + w = −125 + 104 = −21 kJ
31.
a. ΔE = q + w = −47 kJ + 88 kJ = 41 kJ
b. ΔE = 82 − 47 = 35 kJ
c. ΔE = 47 + 0 = 47 kJ
d. When the surroundings do work on the system, w > 0. This is the case for a.
32.
Step 1: ΔE1 = q + w = 72 J + 35 J = 107 J; step 2: ΔE2 = 35 J − 72 J = −37 J
ΔEoverall = ΔE1 + ΔE2 = 107 J − 37 J = 70. J
33.
ΔE = q + w; work is done by the system on the surroundings in a gas expansion; w is
negative.
300. J = q − 75 J, q = 375 J of heat transferred to the system
188
34.
CHAPTER 6
THERMOCHEMISTRY
a. ΔE = q + w = −23 J + 100. J = 77 J
b. w = −PΔV = −1.90 atm(2.80 L − 8.30 L) = 10.5 L atm ×
101.3 J
= 1060 J
L atm
ΔE = q + w = 350. J + 1060 = 1410 J
c. w = −PΔV = −1.00 atm(29.1 L−11.2 L) = −17.9 L atm ×
101.3 J
= −1810 J
L atm
ΔE = q + w = 1037 J − 1810 J = −770 J
35.
w = −PΔV; we need the final volume of the gas. Because T and n are constant, P1V1 = P2V2.
V2 =
V1 P1
P2
=
10.0 L(15.0 atm)
= 75.0 L
2.00 atm
w = −PΔV = −2.00 atm(75.0 L − 10.0 L) = −130. L atm ×
101.3 J
1 kJ
×
L atm
1000 J
= −13.2 kJ = work
36.
w = −210. J = −PΔV, −210 J = −P(25 L − 10. L), P = 14 atm
37.
In this problem, q = w = −950. J.
−950. J ×
1 L atm
= −9.38 L atm of work done by the gases
101.3 J
w = −PΔV, −9.38 L atm =
38.
− 650.
atm × (Vf − 0.040 L), Vf − 0.040 = 11.0 L, Vf = 11.0 L
760
ΔE = q + w, −102.5 J = 52.5 J + w, w = −155.0 J ×
1 L atm
= −1.530 L atm
101.3 J
w = −PΔV, −1.530 L atm = −0.500 atm × ΔV, ΔV = 3.06 L
ΔV = Vf – Vi, 3.06 L = 58.0 L − Vi, Vi = 54.9 L = initial volume
39.
q = molar heat capacity × mol × ΔT =
20.8 J
o
× 39.1 mol × (38.0 − 0.0)°C = 30,900 J
C mol
= 30.9 kJ
w = −PΔV = −1.00 atm × (998 L − 876 L) = −122 L atm ×
ΔE = q + w = 30.9 kJ + (−12.4 kJ) = 18.5 kJ
101.3 J
= −12,400 J = −12.4 kJ
L atm
CHAPTER
40.
6
THERMOCHEMISTRY
189
H2O(g) → H2O(l); ΔE = q + w; q = −40.66 kJ; w = −PΔV
Volume of 1 mol H2O(l) = 1.000 mol H2O(l) ×
18.02 g
1 cm 3
×
= 18.1 cm3 = 18.1 mL
mol
0.996 g
w = −PΔV = −1.00 atm × (0.0181 L − 30.6 L) = 30.6 L atm ×
101.3 J
= 3.10 × 103 J
L atm
= 3.10 kJ
ΔE = q + w = −40.66 kJ + 3.10 kJ = −37.56 kJ
Properties of Enthalpy
41.
This is an endothermic reaction, so heat must be absorbed in order to convert reactants into
products. The high-temperature environment of internal combustion engines provides the
heat.
42.
One should try to cool the reaction mixture or provide some means of removing heat because
the reaction is very exothermic (heat is released). The H2SO4(aq) will get very hot and
possibly boil unless cooling is provided.
43.
a. Heat is absorbed from the water (it gets colder) as KBr dissolves, so this is an
endothermic process.
b. Heat is released as CH4 is burned, so this is an exothermic process.
c. Heat is released to the water (it gets hot) as H2SO4 is added, so this is an exothermic
process.
d. Heat must be added (absorbed) to boil water, so this is an endothermic process.
44.
a. The combustion of gasoline releases heat, so this is an exothermic process.
b. H2O(g) → H2O(l); heat is released when water vapor condenses, so this is an exothermic
process.
c. To convert a solid to a gas, heat must be absorbed, so this is an endothermic process.
d. Heat must be added (absorbed) in order to break a bond, so this is an endothermic
process.
45.
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) ΔH = −1652 kJ; note that 1652 kJ of heat is released when
4 mol Fe reacts with 3 mol O2 to produce 2 mol Fe2O3.
a. 4.00 mol Fe ×
− 1652 kJ
= −1650 kJ; 1650 kJ of heat released
4 mol Fe
b. 1.00 mol Fe2O3 ×
− 1652 kJ
= −826 kJ; 826 kJ of heat released
2 mol Fe 2 O 3
190
CHAPTER 6
THERMOCHEMISTRY
c. 1.00 g Fe ×
1 mol Fe
− 1652 kJ
= −7.39 kJ; 7.39 kJ of heat released
×
55.85 g
4 mol Fe
d. 10.0 g Fe ×
1 mol Fe
− 1652 kJ
= −73.9 kJ
×
55.85 g Fe
4 mol Fe
1 mol O 2
− 1652 kJ
= −34.4 kJ
×
32.00 g O 2
3 mol O 2
2.00 g O2 ×
Because 2.00 g O2 releases the smaller quantity of heat, O2 is the limiting reactant and
34.4 kJ of heat can be released from this mixture.
46.
a. 1.00 mol H2O ×
− 572 kJ
= −286 kJ; 286 kJ of heat released
2 mol H 2 O
b. 4.03 g H2 ×
1 mol H 2
− 572 kJ
= −572 kJ; 572 kJ of heat released
×
2 mol H 2
2.016 g H 2
c. 186 g O2 ×
− 572 kJ
1 mol O 2
= −3320 kJ; 3320 kJ of heat released
×
32.00 g O 2
mol O 2
d.
n H2 =
1.0 atm × 2.0 × 10 8 L
PV
=
= 8.2 × 106 mol H2
0.08206 L atm
RT
× 298 K
K mol
8.2 × 106 mol H2 ×
47.
− 572 kJ
= −2.3 × 109 kJ; 2.3 × 109 kJ of heat released
2 mol H 2
From Example 6.3, q = 1.3 × 108 J. Because the heat transfer process is only 60.%
100. J
efficient, the total energy required is 1.3 × 108 J ×
= 2.2 × 108 J.
60. J
Mass C3H8 = 2.2 × 108 J ×
48.
a. 1.00 g CH4 ×
b. n =
PV
=
RT
2221 × 10 J
3
×
44.09 g C3 H 8
= 4.4 × 103 g C3H8
mol C3 H 8
1 mol CH 4
− 891 kJ
= −55.5 kJ
×
mol CH 4
16.04 g CH 4
1 atm
× 1.00 × 10 3 L
760 torr
= 39.8 mol CH4
0.08206 L atm
× 298 K
K mol
740. torr ×
39.8 mol CH4 ×
49.
1 mol C3 H 8
− 891 kJ
= −3.55 × 104 kJ
mol CH 4
When a liquid is converted into gas, there is an increase in volume. The 2.5 kJ/mol quantity is
the work done by the vaporization process in pushing back the atmosphere.
CHAPTER
50.
6
THERMOCHEMISTRY
191
∆H = ∆E + P∆V; from this equation, ∆H > ∆E when ∆V > 0, ∆H < ∆E when ∆V < 0, and ∆H
= ∆E when ∆V = 0. Concentrate on the moles of gaseous products versus the moles of
gaseous reactants to predict ∆V for a reaction.
a. There are 2 moles of gaseous reactants converting to 2 moles of gaseous products, so
∆V = 0. For this reaction, ∆H = ∆E.
b. There are 4 moles of gaseous reactants converting to 2 moles of gaseous products, so
∆V < 0 and ∆H < ∆E.
c. There are 9 moles of gaseous reactants converting to 10 moles of gaseous products, so
∆V > 0 and ∆H > ∆E.
Calorimetry and Heat Capacity
51.
Specific heat capacity is defined as the amount of heat necessary to raise the temperature of
one gram of substance by one degree Celsius. Therefore, H2O(l) with the largest heat
capacity value requires the largest amount of heat for this process. The amount of heat for
H2O(l) is:
4.18 J
energy = s × m × ΔT = o
× 25.0 g × (37.0°C − 15.0°C) = 2.30 × 103 J
Cg
The largest temperature change when a certain amount of energy is added to a certain mass of
substance will occur for the substance with the smallest specific heat capacity. This is Hg(l),
and the temperature change for this process is:
1000 J
10.7 kJ ×
energy
kJ = 140°C
ΔT =
=
0.14 J
s×m
× 550. g
o
Cg
52.
a. s = specific heat capacity =
Energy = s × m × ΔT =
b. Molar heat capacity =
c. 1250 J =
53.
0.24 J
o
0.24 J
o
Cg
0.24 J
o
0.24 J
since ΔT(K) = ΔT(°C)
Kg
× 150.0 g × (298 K − 273 K) = 9.0 × 102 J
Cg
0.24 J
o
=
Cg
×
107.9 g Ag
=
mol Ag
× m × (15.2°C − 12.0°C), m =
Cg
s = specific heat capacity =
26 J
o
C mol
1250
= 1.6 × 103 g Ag
0.24 × 3.2
q
133 J
= 0.890 J/°C•g
=
m × ∆T
5.00 g × (55.1 − 25.2) o C
From Table 6.1, the substance is solid aluminum.
192
54.
CHAPTER 6
s=
585 J
125.6 g × (53.5 − 20.0) o C
Molar heat capacity =
55.
= 0.139 J/°C•g
0.139 J
o
THERMOCHEMISTRY
Cg
×
200.6 g
27.9 J
= o
mol Hg
C mol
| Heat loss by hot water | = | heat gain by cooler water |
The magnitudes of heat loss and heat gain are equal in calorimetry problems. The only
difference is the sign (positive or negative). To avoid sign errors, keep all quantities positive
and, if necessary, deduce the correct signs at the end of the problem. Water has a specific
heat capacity = s = 4.18 J/°C•g = 4.18 J/K•g (ΔT in °C = ΔT in K).
Heat loss by hot water = s × m × ΔT =
Heat gain by cooler water =
4.18 J
× 50.0 g × (330. K − Tf)
Kg
4.18 J
× 30.0 g × (Tf − 280. K); heat loss = heat gain, so:
Kg
209 J
125 J
× (330. K − Tf) =
× (Tf − 280. K)
K
K
6.90 × 104 − 209Tf = 125Tf − 3.50 × 104, 334Tf = 1.040 × 105, Tf = 311 K
Note that the final temperature is closer to the temperature of the more massive hot water,
which is as it should be.
56.
Heat loss by hot water = heat gain by cold water; keeping all quantities positive helps to
avoid sign errors:
4.18 J
o
Cg
mhot =
57.
× mhot × (55.0°C − 37.0°C) =
4.18 J
o
× 90.0 g × (37.0°C − 22.0°C)
Cg
90.0 g × 15.0o C
= 75.0 g hot water needed
18.0o C
Heat loss by Al + heat loss by Fe = heat gain by water; keeping all quantities positive to
avoid sign error:
0.45 J
0.89 J
× 5.00 g Al × (100.0°C − Tf) + o
× 10.00 g Fe × (100.0 − Tf)
o
Cg
Cg
4.18 J
× 97.3 g H2O × (Tf − 22.0°C)
= o
Cg
4.5(100.0 − Tf) + 4.5(100.0 − Tf) = 407(Tf − 22.0), 450 − (4.5)Tf + 450 − (4.5)Tf
= 407Tf − 8950
416Tf = 9850, Tf = 23.7°C
CHAPTER
58.
6
THERMOCHEMISTRY
Heat released to water = 5.0 g H2 ×
Heat gain by water = 1.10 × 103 J =
193
120. J
50. J
+ 10. g methane ×
= 1.10 × 103 J
g H2
g methane
4.18 J
o
× 50.0 g × ∆T
Cg
∆T = 5.26°C, 5.26°C = Tf − 25.0°C, Tf = 30.3°C
59.
Heat gain by water = heat loss by metal = s × m × ΔT, where s = specific heat capacity.
Heat gain =
4.18 J
o
× 150.0 g × (18.3°C − 15.0°C) = 2100 J
Cg
A common error in calorimetry problems is sign errors. Keeping all quantities positive helps
to eliminate sign errors.
Heat loss = 2100 J = s × 150.0 g × (75.0°C − 18.3°C), s =
60.
2100 J
150.0 g × 56.7 o C
= 0.25 J/°C•g
Heat gain by water = heat loss by Cu; keeping all quantities positive helps to avoid sign
errors:
4.18 J
o
× mass × (24.9°C − 22.3°C) =
Cg
0.20 J
o
× 110. g Cu × (82.4°C − 24.9°C)
Cg
11(mass) = 1300, mass = 120 g H2O
61.
50.0 × 10−3 L × 0.100 mol/L = 5.00 × 10−3 mol of both AgNO3 and HCl are reacted. Thus
5.00 × 10−3 mol of AgCl will be produced because there is a 1 : 1 mole ratio between
reactants.
Heat lost by chemicals = heat gained by solution
Heat gain =
4.18 J
× 100.0 g × (23.40 − 22.60)°C = 330 J
o
Cg
Heat loss = 330 J; this is the heat evolved (exothermic reaction) when 5.00 × 10−3 mol of
AgCl is produced. So q = −330 J and ΔH (heat per mol AgCl formed) is negative with a
value of:
ΔH =
− 330 J
1 kJ
= −66 kJ/mol
×
−3
5.00 × 10 mol 1000 J
Note: Sign errors are common with calorimetry problems. However, the correct sign for ΔH
can be determined easily from the ΔT data; i.e., if ΔT of the solution increases, then the
reaction is exothermic because heat was released, and if ΔT of the solution decreases, then
the reaction is endothermic because the reaction absorbed heat from the water. For
calorimetry problems, keep all quantities positive until the end of the calculation and then
decide the sign for ΔH. This will help eliminate sign errors.
194
62.
CHAPTER 6
THERMOCHEMISTRY
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
We have a stoichiometric mixture. All of the NaOH and HCl will react.
0.10 L ×
1.0 mol
= 0.10 mol of HCl is neutralized by 0.10 mol NaOH.
L
Heat lost by chemicals = heat gained by solution
Volume of solution = 100.0 + 100.0 = 200.0 mL
Heat gain =
1.0 g 

3
×  200.0 mL ×
 × (31.3 – 24.6)°C = 5.6 × 10 J = 5.6 kJ
mL 
Cg 
4.18 J
o
Heat loss = 5.6 kJ; this is the heat released by the neutralization of 0.10 mol HCl. Because
the temperature increased, the sign for ΔH must be negative, i.e., the reaction is exothermic.
For calorimetry problems, keep all quantities positive until the end of the calculation and then
decide the sign for ΔH. This will help eliminate sign errors.
ΔH =
63.
− 5.6 kJ
= −56 kJ/mol
0.10 mol
Heat lost by solution = heat gained by KBr; mass of solution = 125 g + 10.5 g = 136 g
Note: Sign errors are common with calorimetry problems. However, the correct
ΔH can easily be obtained from the ΔT data. When working calorimetry problems,
quantities positive (ignore signs). When finished, deduce the correct sign for ΔH.
problem, T decreases as KBr dissolves, so ΔH is positive; the dissolution of
endothermic (absorbs heat).
Heat lost by solution =
ΔH in units of J/g =
4.18 J
× 136 g × (24.2°C − 21.1°C) = 1800 J = heat gained by KBr
o
Cg
1800 J
= 170 J/g
10.5 g KBr
ΔH in units of kJ/mol =
64.
sign for
keep all
For this
KBr is
170 J
119.0 g KBr
1 kJ
= 20. kJ/mol
×
×
g KBr
mol KBr
1000 J
NH4NO3(s) → NH4+(aq) + NO3−(aq) ΔH = ?; mass of solution = 75.0 g + 1.60 g = 76.6 g
Heat lost by solution = heat gained as NH4NO3 dissolves. To help eliminate sign errors, we
will keep all quantities positive (q and ΔT) and then deduce the correct sign for ΔH at the end
of the problem. Here, because temperature decreases as NH4NO3 dissolves, heat is absorbed
as NH4NO3 dissolves, so this is an endothermic process (ΔH is positive).
Heat lost by solution =
4.18 J
× 76.6 g × (25.00 − 23.34)°C = 532 J = heat gained as
o
Cg
NH4NO3 dissolves
CHAPTER
ΔH =
65.
6
THERMOCHEMISTRY
195
80.05 g NH 4 NO 3
532 J
1 kJ
= 26.6 kJ/mol NH4NO3 dissolving
×
×
1.60 g NH 4 NO 3
mol NH 4 NO 3
1000 J
Because ΔH is exothermic, the temperature of the solution will increase as CaCl2(s)
dissolves. Keeping all quantities positive:
heat loss as CaCl2 dissolves = 11.0 g CaCl2 ×
heat gained by solution = 8.08 × 103 J =
1 mol CaCl 2
81.5 kJ
= 8.08 kJ
×
110.98 g CaCl 2
mol CaCl 2
4.18 J
× (125 + 11.0) g × (Tf − 25.0°C)
o
Cg
8.08 × 103
Tf − 25.0°C =
= 14.2°C, Tf = 14.2°C + 25.0°C = 39.2°C
4.18 × 136
66.
0.1000 L ×
0.500 mol HCl 118 kJ heat released
×
= 2.95 kJ of heat released if HCl limiting
L
2 mol HCl
0.3000 L ×
118 kJ heat released
0.100 mol Ba (OH) 2
×
= 3.54 kJ heat released if
mol Ba (OH) 2
L
Ba(OH)2 limiting
Because the HCl reagent produces the smaller amount of heat released, HCl is limiting and
2.95 kJ of heat are released by this reaction.
Heat gained by solution = 2.95 × 103 J =
4.18 J
× 400.0 g × ΔT
o
Cg
ΔT = 1.76°C = Tf − Ti = Tf − 25.0°C, Tf = 26.8°C
67.
a. Heat gain by calorimeter = heat loss by CH4 = 6.79 g CH4 ×
Heat capacity of calorimeter =
1 mol CH 4
802 kJ
×
16.04 g
mol
= 340. kJ
340. kJ
= 31.5 kJ/°C
10.8 o C
b. Heat loss by C2H2 = heat gain by calorimeter = 16.9°C ×
31. 5 kJ
= 532 kJ
o
C
A bomb calorimeter is at constant volume, so the heat released/gained = qV = ∆E:
ΔEcomb =
68.
26.04 g
− 532 kJ
= −1.10 × 103 kJ/mol
×
mol C 2 H 2
12.6 g C 2 H 2
First, we need to get the heat capacity of the calorimeter from the combustion of benzoic
acid. Heat lost by combustion = heat gained by calorimeter.
Heat loss = 0.1584 g ×
26.42 kJ
= 4.185 kJ
g
196
CHAPTER 6
Heat gain = 4.185 kJ = Ccal × ΔT, Ccal =
4.185 kJ
2.54 o C
THERMOCHEMISTRY
= 1.65 kJ/°C
Now we can calculate the heat of combustion of vanillin. Heat loss = heat gain.
Heat gain by calorimeter =
1.65 kJ
o
× 3.25°C = 5.36 kJ
C
Heat loss = 5.36 kJ, which is the heat evolved by combustion of the vanillin.
∆Ecomb =
− 5.36 kJ
− 25.2 kJ 152.14 g
= −25.2 kJ/g; ∆Ecomb =
= −3830 kJ/mol
×
0.2130 g
g
mol
Hess's Law
69.
Information given:
C(s) + O2(g) → CO2(g)
CO(g) + 1/2 O2(g) → CO2(g)
ΔH = −393.7 kJ
ΔH = −283.3 kJ
Using Hess’s law:
2 C(s) + 2 O2(g) → 2 CO2(g)
2 CO2(g) → 2 CO(g) + O2(g)
2 C(s) + O2(g) → 2 CO(g)
ΔH1 = 2(−393.7 kJ)
ΔH2 = −2(−283.3 kJ)
ΔH = ΔH1 + ΔH2 = −220.8 kJ
Note: When an equation is reversed, the sign on ΔH is reversed. When the coefficients in a
balanced equation are multiplied by an integer, then the value of ΔH is multiplied by the
same integer.
70.
Given:
C4H4(g) + 5 O2(g) → 4 CO2(g) + 2 H2O(l)
C4H8(g) + 6 O2(g) → 4 CO2(g) + 4 H2O(l)
H2(g) + 1/2 O2(g) → H2O(l)
ΔHcomb = −2341 kJ
ΔHcomb = −2755 kJ
ΔHcomb = −286 kJ
By convention, H2O(l) is produced when enthalpies of combustion are given, and because
per-mole quantities are given, the combustion reaction refers to 1 mole of that quantity
reacting with O2(g).
Using Hess’s law to solve:
C4H4(g) + 5 O2(g) → 4 CO2(g) + 2 H2O(l)
4 CO2(g) + 4 H2O(l) → C4H8(g) + 6 O2(g)
2 H2(g) + O2(g) → 2 H2O(l)
C4H4(g) + 2 H2(g) → C4H8(g)
ΔH1 = −2341 kJ
ΔH2 = − (−2755 kJ)
ΔH3 = 2(−286 kJ)
ΔH = ΔH1 + ΔH2 + ΔH3 = −158 kJ
CHAPTER
71.
6
THERMOCHEMISTRY
2 N2(g) + 6 H2(g) → 4 NH3(g)
6 H2O(g) → 6 H2(g) + 3 O2(g)
2 N2(g) + 6 H2O(g) → 3 O2(g) + 4 NH3(g)
197
ΔH = −2(92 kJ)
ΔH = −3(−484 kJ)
ΔH = 1268 kJ
No, because the reaction is very endothermic (requires a lot of heat to react), it would not be a
practical way of making ammonia because of the high energy costs required.
72.
ΔH = 1/2(167.4 kJ)
ΔH = −1/2(341.4 kJ)
ΔH = 1/2(−43.4 kJ)
ClF + 1/2 O2 → 1/2 Cl2O + 1/2 F2O
1/2 Cl2O + 3/2 F2O → ClF3 + O2
F2 + 1/2 O2 → F2O
ΔH = −108.7 kJ
ClF(g) + F2(g) → ClF3
73.
NO + O3 → NO2 + O2
3/2 O2 → O3
O → 1/2 O2
NO(g) + O(g) → NO2(g)
74.
ΔH = −199 kJ
ΔH = −1/2(−427 kJ)
ΔH = −1/2(495 kJ)
ΔH = −233 kJ
We want ΔH for N2H4(l) + O2(g) → N2(g) + 2 H2O(l). It will be easier to calculate ΔH for
the combustion of four moles of N2H4 because we will avoid fractions.
ΔH = 9(−286 kJ)
ΔH = −3(−317 kJ)
ΔH = −1010. kJ
ΔH = − (−143 kJ)
9 H2 + 9/2 O2 → 9 H2O
3 N2H4 + 3 H2O → 3 N2O + 9 H2
2 NH3 + 3 N2O → 4 N2 + 3 H2O
N2H4 + H2O → 2 NH3 + 1/2 O2
ΔH = −2490. kJ
4 N2H4(l) + 4 O2(g) → 4 N2(g) + 8 H2O(l)
For N2H4(l) + O2(g) → N2(g) + 2 H2O(l)
ΔH =
− 2490. kJ
= −623 kJ
4
Note: By the significant figure rules, we could report this answer to four significant figures.
However, because the ΔH values given in the problem are only known to ±1 kJ, our final
answer will at best be ±1 kJ.
75.
CaC2 →
CaO + H2O →
2 CO2 + H2O →
Ca + 1/2 O2 →
2 C + 2 O2 →
Ca + 2 C
Ca(OH)2
C2H2 + 5/2 O2
CaO
2 CO2
CaC2(s) + 2 H2O(l) → Ca(OH)2(aq) + C2H2(g)
ΔH = − (−62.8 kJ)
ΔH = −653.1 kJ
ΔH = − (−1300. kJ)
ΔH = −635.5 kJ
ΔH = 2(−393.5 kJ)
ΔH = −713 kJ
198
CHAPTER 6
P4O10 → P4 + 5 O2
10 PCl3 + 5 O2 → 10 Cl3PO
6 PCl5 → 6 PCl3 + 6 Cl2
P4 + 6 Cl2 → 4 PCl3
76.
THERMOCHEMISTRY
ΔH = − (−2967.3 kJ)
ΔH = 10(−285.7 kJ)
ΔH = −6(−84.2 kJ)
ΔH = −1225.6
P4O10(s) + 6 PCl5(g) → 10 Cl3PO(g)
ΔH = −610.1 kJ
Standard Enthalpies of Formation
77.
The change in enthalpy that accompanies the formation of 1 mole of a compound from its
elements, with all substances in their standard states, is the standard enthalpy of formation for
a compound. The reactions that refer to ∆H of are:
Na(s) + 1/2 Cl2(g) → NaCl(s); H2(g) + 1/2 O2(g) → H2O(l)
6 C(graphite, s) + 6 H2(g) + 3 O2(g) → C6H12O6(s)
Pb(s) + S(rhombic, s) + 2 O2(g) → PbSO4(s)
78.
a. Aluminum oxide = Al2O3; 2 Al(s) + 3/2 O2(g) → Al2O3(s)
b. C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(l)
c. NaOH(aq) + HCl(aq) → H2O(l) + NaCl(aq)
d. 2 C(graphite, s) + 3/2 H2(g) + 1/2 Cl2(g) → C2H3Cl(g)
e. C6H6(l) + 15/2 O2(g) → 6 CO2(g) + 3 H2O(l)
Note: ΔHcomb values assume 1 mole of compound combusted.
f.
79.
NH4Br(s) → NH4+(aq) + Br−(aq)
In general, ΔH° = ∑ np ΔH f , products − ∑ nr ΔH f , reactants , and all elements in their standard
state have ΔH f = 0 by definition.
a. The balanced equation is 2 NH3(g) + 3 O2(g) + 2 CH4(g) → 2 HCN(g) + 6 H2O(g).
ΔH° = (2 mol HCN × ΔH f , HCN + 6 mol H2O(g) × ΔH f , H 2O )
− (2 mol NH3 × ΔH f , NH 3 + 2 mol CH4 × ∆H f , CH )
4
ΔH° = [2(135.1) + 6(−242)] − [2(−46) + 2(−75)] = −940. kJ
CHAPTER
6
THERMOCHEMISTRY
199
b. Ca3(PO4)2(s) + 3 H2SO4(l) → 3 CaSO4(s) + 2 H3PO4(l)

 − 1267 kJ  
 − 1433 kJ 
ΔH° = 3 mol CaSO 4 (s)

 + 2 mol H 3 PO 4 (l)
 mol  
 mol 


 − 814 kJ  
 − 4126 kJ 
− 1 mol Ca 3 (PO 4 ) 2 (s)
 + 3 mol H 2SO 4 (l)

 mol 
 mol  

ΔH° = −6833 kJ − (−6568 kJ) = −265 kJ
c. NH3(g) + HCl(g) → NH4Cl(s)
ΔH° = (1 mol NH4Cl × ΔH f , NH 4Cl ) − (1 mol NH3 × ΔH f , NH 3 + 1 mol HCl × ΔH f , HCl )

 − 314 kJ  
 − 46 kJ 
 − 92 kJ 
ΔH° = 1 mol 
 − 1 mol 
 + 1 mol 

 mol  
 mol 
 mol 

ΔH° = −314 kJ + 138 kJ = −176 kJ
80.
a. The balanced equation is C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(g).

 − 393.5 kJ 
 − 242 kJ  
 − 278 kJ 
ΔH° = 2 mol 
 + 3 mol 
 − 1 mol 

 mol 
 mol  
 mol 

ΔH° = −1513 kJ − (−278 kJ) = −1235 kJ
b. SiCl4(l) + 2 H2O(l) → SiO2(s) + 4 HCl(aq)
Because HCl(aq) is H+(aq) + Cl−(aq), ΔH f = 0 − 167 = −167 kJ/mol.


 − 167 kJ 
 − 911 kJ 
 − 687 kJ 
 − 286 kJ 
ΔH° = 4 mol 
 + 1 mol 
 − 1 mol 
 + 2 mol 

 mol 
 mol 
 mol 
 mol 


ΔH° = −1579 kJ − (−1259 kJ) = −320. kJ
c. MgO(s) + H2O(l) → Mg(OH)2(s)

 − 286 kJ 
 − 602 kJ 
 − 925 kJ  
ΔH° = 1 mol 

 + 1 mol
 − 1 mol 
 mol 
 mol 
 mol  

ΔH° = −925 kJ − (−888 kJ) = −37 kJ
81.
a. 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g); ΔH° = ∑ np ΔH f , products − ∑ nr ΔH f , reactants

 90. kJ 
 − 242 kJ  
 − 46 kJ 
ΔH° = 4 mol 
 + 6 mol 
 − 4 mol 
 = −908 kJ
 mol 
 mol  
 mol 

200
CHAPTER 6
THERMOCHEMISTRY
2 NO(g) + O2(g) → 2 NO2(g)

 90. kJ  
 34 kJ  
ΔH° = 2 mol 
  = −112 kJ
 − 2 mol 
 mol  
 mol  

3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g)

 − 286 kJ 
 − 207 kJ 
 90. kJ   
 34 kJ 
ΔH° = 2 mol 

 + 1 mol 
  − 3 mol 
 + 1 mol 
 mol 
 mol 
 mol   
 mol 

−140. kJ
Note: All ΔH f values are assumed ±1 kJ.
b. 12 NH3(g) + 15 O2(g) → 12 NO(g) + 18 H2O(g)
12 NO(g) + 6 O2(g) → 12 NO2(g)
12 NO2(g) + 4 H2O(l) → 8 HNO3(aq) + 4 NO(g)
4 H2O(g) → 4 H2O(l)
12 NH3(g) + 21 O2(g) → 8 HNO3(aq) + 4 NO(g) + 14 H2O(g)
The overall reaction is exothermic because each step is exothermic.
82.
 − 416 kJ 
4 Na(s) + O2(g) → 2 Na2O(s) ΔH° = 2 mol 
 = −832 kJ
 mol 
2 Na(s) + 2 H2O(l) → 2 NaOH(aq) + H2(g)

 − 470. kJ   
 − 286 kJ 
ΔH° = 2 mol 
  − 2 mol 
 = −368 kJ
 mol   
 mol 

2Na(s) + CO2(g) → Na2O(s) + CO(g)

 − 393.5 kJ 
 − 110.5 kJ  
 − 416 kJ 
ΔH° = 1 mol 
 = −133 kJ
 − 1 mol 
 + 1 mol 
 mol 
 mol  
 mol 

In Reactions 2 and 3, sodium metal reacts with the "extinguishing agent." Both reactions are
exothermic, and each reaction produces a flammable gas, H2 and CO, respectively.
83.
3 Al(s) + 3 NH4ClO4(s) → Al2O3(s) + AlCl3(s) + 3 NO(g) + 6 H2O(g)

 − 1676 kJ 
 − 704 kJ 
 − 242 kJ 
 90. kJ 
ΔH° = 6 mol 

 + 1 mol 
 + 3 mol 
 + 1 mol 
 mol 
 mol 
 mol 
 mol 


 − 295 kJ 
− 3 mol 
 = −2677 kJ
 mol 

CHAPTER
84.
6
THERMOCHEMISTRY
201
5 N2O4(l) + 4 N2H3CH3(l) → 12 H2O(g) + 9 N2(g) + 4 CO2(g)

 − 242 kJ 
 − 393.5 kJ  
ΔH° = 12 mol 
 + 4 mol 

 mol 
 mol  


 − 20. kJ 
− 5 mol 
 + 4 mol
 mol 

85.
2 ClF3(g) + 2 NH3(g) → N2(g) + 6 HF(g) + Cl2(g)
ΔH° = (6 ΔH of , HF ) − (2 ∆H of ,
ClF3
+ 2∆H of ,
 54 kJ 
 = −4594 kJ

 mol 
ΔH° = −1196 kJ
NH 3 )
 − 271 kJ 
 − 46 kJ 
o
−1196 kJ = 6 mol 
 − 2 ΔH f , ClF3 − 2 mol 

mol


 mol 
−1196 kJ = −1626 kJ − 2 ΔH of , ClF3 + 92 kJ, ΔH of , ClF3 =
86.
C2H4(g) + 3 O2(g) → 2 CO2(g) + 2 H2O(l)
(−1626 + 92 + 1196) kJ − 169 kJ
=
mol
2 mol
ΔH° = −1411.1 kJ
ΔH° = −1411.1 kJ = 2(−393.5) kJ + 2(−285.8) kJ − ΔH f , C 2 H 4
−1411.1 kJ = −1358.6 kJ − ΔH f , C 2 H 4 , ΔH f , C 2 H 4 = 52.5 kJ/mol
Energy Consumption and Sources
87.
C(s) + H2O(g) → H2(g) + CO(g) ΔH° = −110.5 kJ − (−242 kJ) = 132 kJ
88.
CO(g) + 2 H2(g) → CH3OH(l)
89.
C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(l)
ΔH° = −239 kJ − (−110.5 kJ) = −129 kJ
ΔH° = [2(−393.5 kJ) + 3(−286 kJ)] − (−278 kJ) = −1367 kJ/mol ethanol
1 mol
− 1367 kJ
= −29.67 kJ/g
×
mol
46.07 g
90.
CH3OH(l) + 3/2 O2(g) → CO2(g) + 2 H2O(l)
ΔH° = [−393.5 kJ + 2(−286 kJ)] − (−239 kJ) = −727 kJ/mol CH3OH
− 727 kJ
1 mol
= −22.7 kJ/g versus −29.67 kJ/g for ethanol (from Exercise 89)
×
mol
32.04 g
Ethanol has a slightly higher fuel value per gram than methanol.
202
91.
CHAPTER 6
THERMOCHEMISTRY
C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(l)
ΔH° = [3(−393.5 kJ) + 4(−286 kJ)] − (−104 kJ) = −2221 kJ/mol C3H8
− 50.37 kJ
− 2221 kJ
1 mol
=
versus −47.7 kJ/g for octane (Example 6.11)
×
g
mol
44.09 g
The fuel values are very close. An advantage of propane is that it burns more cleanly. The
boiling point of propane is −42°C. Thus it is more difficult to store propane, and there are
extra safety hazards associated with using high-pressure compressed-gas tanks.
92.
1 mole of C2H2(g) and 1 mole of C4H10(g) have equivalent volumes at the same T and P.
enthalpy of combustion per mol of C 2 H 2
Enthalpy of combustion per volume of C 2 H 2
=
Enthalpy of combustion per volume of C 4 H10
enthalpy of combustion per mol of C 4 H 10
Enthalpy of combustion per volume of C 2 H 2
=
Enthalpy of combustion per volume of C 4 H10
− 49.9 kJ 26.04 g C 2 H 2
×
g C2H 2
mol C 2 H 2
= 0.452
− 49.5 kJ 58.12 g C 4 H10
×
g C 4 H10
mol C 4 H10
More than twice the volume of acetylene is needed to furnish the same energy as a given
volume of butane.
93.
The molar volume of a gas at STP is 22.42 L (from Chapter 5).
4.19 × 106 kJ ×
94.
1 mol CH 4
22.42 L CH 4
= 1.05 × 105 L CH4
×
891 kJ
mol CH 4
Mass of H2O = 1.00 gal ×
3.785 L 1000 mL 1.00 g
= 3790 g H2O
×
×
gal
L
mL
Energy required (theoretical) = s × m × ΔT =
4.18 J
× 3790 g × 10.0 °C = 1.58 × 105 J
o
Cg
For an actual (80.0% efficient) process, more than this quantity of energy is needed since heat
is always lost in any transfer of energy. The energy required is:
1.58 × 105 J ×
100. J
= 1.98 × 105 J
80.0 J
Mass of C2H2 = 1.98 × 105 J ×
1 mol C 2 H 2
26.04 g C 2 H 2
= 3.97 g C2H2
×
3
mol C 2 H 2
1300. × 10 J
Additional Exercises
95.
2.0 h ×
1 mol H 2 O
18.02 g H 2 O
5500 kJ
= 4900 g = 4.9 kg H2O
×
×
h
40.6 kJ
mol
CHAPTER
96.
6
203
From the problem, walking 4.0 miles consumes 400 kcal of energy.
1 lb fat ×
97.
THERMOCHEMISTRY
454 g
7.7 kcal
4 mi
1h
= 8.7 h = 9 h
×
×
×
lb
g
400 kcal 4 mi
a. 2 SO2(g) + O2(g) → 2 SO3(g); w = −PΔV; because the volume of the piston apparatus
decreased as reactants were converted to products (∆V < 0), w is positive (w > 0).
b. COCl2(g) → CO(g) + Cl2(g); because the volume increased (∆V > 0), w is negative
(w < 0).
c. N2(g) + O2(g) → 2 NO(g); because the volume did not change (∆V = 0), no PV work is
done (w = 0).
In order to predict the sign of w for a reaction, compare the coefficients of all the product
gases in the balanced equation to the coefficients of all the reactant gases. When a balanced
reaction has more moles of product gases than moles of reactant gases (as in b), the reaction
will expand in volume (ΔV positive), and the system does work on the surroundings. When
a balanced reaction has a decrease in the moles of gas from reactants to products (as in a), the
reaction will contract in volume (ΔV negative), and the surroundings will do compression
work on the system. When there is no change in the moles of gas from reactants to products
(as in c), ΔV = 0 and w = 0.
98.
a. N2(g) + 3 H2(g) → 2 NH3(g); from the balanced equation, 1 molecule of N2 will react
with 3 molecules of H2 to produce 2 molecules of NH3. So the picture after the reaction
should only have 2 molecules of NH3 present. Another important part of your drawing
will be the relative volume of the product container. The volume of a gas is directly
proportional to the number of gas molecules present (at constant T and P). In this
problem, 4 total molecules of gas were present initially (1 N2 + 3 H2). After reaction,
only 2 molecules are present (2 NH3). Because the number of gas molecules decreases
by a factor of 2 (from 4 total to 2 total), the volume of the product gas must decrease by a
factor of 2 as compared to the initial volume of the reactant gases. Summarizing, the
picture of the product container should have 2 molecules of NH3 and should be at a
volume which is one-half the original reactant container volume.
b. w = −P∆V; here the volume decreased, so ∆V is negative. When ∆V is negative, w is
positive. As the reactants were converted to products, a compression occurred which is
associated with work flowing into the system (w is positive).
99.
a. C12H22O11(s) + 12 O2(g) → 12 CO2(g) + 11 H2O(l)
b. A bomb calorimeter is at constant volume, so heat released = qV = ΔE:
ΔE =
− 24.00 kJ 342.30 g
= −5630 kJ/mol C12H22O11
×
mol
1.46 g
c. PV = nRT; at constant P and T, PΔV = RTΔn, where Δn = moles of gaseous products −
moles of gaseous reactants.
At constant P and T: ΔH = ΔE + PΔV = ΔE + RTΔn
For this reaction, Δn = 12 − 12 = 0, so ΔH = ΔE = −5630 kJ/mol.
204
100.
CHAPTER 6
THERMOCHEMISTRY
w = −PΔV; Δn = moles of gaseous products − moles of gaseous reactants. Only gases can do
PV work (we ignore solids and liquids). When a balanced reaction has more moles of
product gases than moles of reactant gases (Δn positive), the reaction will expand in volume
(ΔV positive), and the system will do work on the surroundings. For example, in reaction c,
Δn = 2 − 0 = 2 moles, and this reaction would do expansion work against the surroundings.
When a balanced reaction has a decrease in the moles of gas from reactants to products (Δn
negative), the reaction will contract in volume (ΔV negative), and the surroundings will do
compression work on the system, e.g., reaction a, where Δn = 0 − 1 = −1. When there is no
change in the moles of gas from reactants to products, ΔV = 0 and w = 0, e.g., reaction b,
where Δn = 2 − 2 = 0.
When ΔV > 0 (Δn > 0), then w < 0, and the system does work on the surroundings (c and e).
When ΔV < 0 (Δn < 0), then w > 0, and the surroundings do work on the system (a and d).
When ΔV = 0 (Δn = 0), then w = 0 (b).
101.
ΔEoverall = ΔEstep 1 + ΔEstep 2; this is a cyclic process, which means that the overall initial state
and final state are the same. Because ΔE is a state function, ΔEoverall = 0 and ΔEstep 1 = −
ΔEstep 2.
ΔEstep 1 = q + w = 45 J + (−10. J) = 35 J
ΔEstep 2 = −ΔEstep 1 = −35 J = q + w, −35 J = −60 J + w, w = 25 J
102.
2 K(s) + 2 H2O(l) → 2 KOH(aq) + H2(g)
5.00 g K ×
ΔH° = 2(−481 kJ) − 2(−286 kJ) = −390. kJ
1 mol K
− 390. kJ
= −24.9 kJ
×
2 mol K
39.10 g K
24.9 kJ of heat is released on reaction of 5.00 g K.
24,900 J =
24,900
4.18 J
× (1.00 × 103 g) × ΔT, ΔT =
= 5.96°C
4.18 × 1.00 × 103
g oC
Final temperature = 24.0 + 5.96 = 30.0°C
103.
HNO3(aq) + KOH(aq) → H2O(l) + KNO3(aq)
ΔH = −56 kJ
0.2000 L ×
0.400 mol HNO 3
56 kJ heat released
×
= 4.5 kJ heat released if HNO3 limiting
mol HNO3
L
0.1500 L ×
0.500 mol KOH
56 kJ heat released
×
= 4.2 kJ heat released if KOH limiting
L
mol KOH
Because the KOH reagent produces the smaller quantity of heat released, KOH is limiting
and 4.2 kJ of heat released.
CHAPTER
104.
6
THERMOCHEMISTRY
205
Na2SO4(aq) + Ba(NO3)2(aq) → BaSO4(s) + 2 NaNO3(aq)
ΔH = ?
1.00 L ×
2.00 mol Na 2 SO 4
1 mol BaSO 4
×
= 2.00 mol BaSO4 if Na2SO4 limiting
L
mol Na 2 SO 4
2.00 L ×
0.750 mol Ba ( NO 3 ) 2
1 mol BaSO 4
×
= 1.50 mol BaSO4 if Ba(NO3)2 limiting
L
mol Ba ( NO3 ) 2
The Ba(NO3)2 reagent produces the smaller quantity of product, so Ba(NO3)2 is limiting and
1.50 mol BaSO4 can form.
Heat gain by solution = heat loss by reaction
Mass of solution = 3.00 L ×
Heat gain by solution =
1000 mL
2.00 g
×
= 6.00 × 103 g
mL
L
6.37 J
o
× 6.00 × 103 g × (42.0 − 30.0)°C = 4.59 × 105 J
Cg
Because the solution gained heat, the reaction is exothermic; q = −4.59 × 105 J for the
reaction.
∆H =
− 4.59 × 105 J
= −3.06 × 105 J/mol = −306 kJ/mol
1.50 mol BaSO 4
105.
|qsurr| = |qsolution + qcal|; we normally assume that qcal is zero (no heat gain/loss by the calorimeter). However, if the calorimeter has a nonzero heat capacity, then some of the heat
absorbed by the endothermic reaction came from the calorimeter. If we ignore qcal, then qsurr is
too small, giving a calculated ∆H value that is less positive (smaller) than it should be.
106.
The specific heat of water is 4.18 J/°C•g, which is equal to 4.18 kJ/°C•kg.
We have 1.00 kg of H2O, so: 1.00 kg ×
4.18 kJ
= 4.18 kJ/°C
o
C kg
This is the portion of the heat capacity that can be attributed to H2O.
Total heat capacity = Ccal + C H 2O , Ccal = 10.84 − 4.18 = 6.66 kJ/°C
107.
Heat released = 1.056 g × 26.42 kJ/g = 27.90 kJ = heat gain by water and calorimeter
 4.18 kJ


 6.66 kJ
Heat gain = 27.90 kJ =  o
× 0.987 kg × ΔT  +  o
× ΔT 
C


 C kg

27.90 = (4.13 + 6.66)ΔT = (10.79)ΔT, ΔT = 2.586°C
2.586°C = Tf − 23.32°C, Tf = 25.91°C
206
108.
CHAPTER 6
THERMOCHEMISTRY
For Exercise 83, a mixture of 3 mol Al and 3 mol NH4ClO4 yields 2677 kJ of energy. The
mass of the stoichiometric reactant mixture is:
26.98 g  
117.49 g 

 3 mol ×
 +  3 mol ×
 = 433.41 g
mol  
mol 

For 1.000 kg of fuel: 1.000 × 103 g ×
− 2677 kJ
= −6177 kJ
433.41 g
In Exercise 84, we get 4594 kJ of energy from 5 mol of N2O4 and 4 mol of N2H3CH3. The
46.08 g 
92.02 g  

mass is  5 mol ×
 = 644.42 kJ.
 +  4 mol ×
mol 
mol  

For 1.000 kg of fuel: 1.000 × 103 g ×
− 4594 kJ
= −7129 kJ
644.42 g
Thus we get more energy per kilogram from the N2O4/N2H3CH3 mixture.
109.
110.
1/2 D → 1/2 A + B
1/2 E + F → 1/2 A
1/2 C → 1/2 E + 3/2 D
∆H = −1/6(−403 kJ)
∆H = 1/2(−105.2 kJ)
∆H = 1/2(64.8 kJ)
F + 1/2 C → A + B + D
∆H = 47.0 kJ
To avoid fractions, let's first calculate ΔH for the reaction:
6 FeO(s) + 6 CO(g) → 6 Fe(s) + 6 CO2(g)
6 FeO + 2 CO2 → 2 Fe3O4 + 2 CO
2 Fe3O4 + CO2 → 3 Fe2O3 + CO
3 Fe2O3 + 9 CO → 6 Fe + 9 CO2
6 FeO(s) + 6 CO(g) → 6 Fe(s) + 6 CO2(g)
So for FeO(s) + CO(g) → Fe(s) + CO2(g), ΔH° =
111.
ΔH° = −2(18 kJ)
ΔH° = − (−39 kJ)
ΔH° = 3(−23 kJ)
ΔH° = −66 kJ
− 66 kJ
= −11 kJ.
6
a. ΔH° = 3 mol(227 kJ/mol) − 1 mol(49 kJ/mol) = 632 kJ
b. Because 3 C2H2(g) is higher in energy than C6H6(l), acetylene will release more energy
per gram when burned in air. Note that 3 moles of C2H2 has the same mass as 1 mole of
C6H6.
CHAPTER
6
THERMOCHEMISTRY
ΔH = − (211.3 kJ)
ΔH = 1/2(242.3 kJ)
ΔH = 1/2(151.0 kJ)
ΔH = 1/2(62.8 kJ)
I(g) + Cl(g) → ICl(g)
1/2 Cl2(g) → Cl(g)
1/2 I2(g) → I(g)
1/2 I2(s) → 1/2 I2(g)
112.
ΔH = 16.8 kJ/mol = ΔH of , ICl
1/2 I2(s) + 1/2 Cl2(g) → ICl(g)
113.
207
Heat gained by water = heat lost by nickel = s × m × ΔT, where s = specific heat capacity.
Heat gain =
4.18 J
o
× 150.0 g × (25.0°C − 23.5°C) = 940 J
Cg
A common error in calorimetry problems is sign errors. Keeping all quantities positive helps
to eliminate sign errors.
Heat loss = 940 J =
114.
0.444 J
o
× mass × (99.8 − 25.0) °C, mass =
Cg
Heat gain by calorimeter =
940
= 28 g
0.444 × 74.8
1.56 kJ
× 3.2°C = 5.0 kJ = heat loss by quinine
o
C
Heat loss = 5.0 kJ, which is the heat evolved (exothermic reaction) by the combustion of
0.1964 g of quinone. Because we are at constant volume, qV = ∆E.
ΔEcomb =
115.
− 5.0 kJ
= −25 kJ/g;
0.1964 g
ΔEcomb =
− 25 kJ 108.09 g
= −2700 kJ/mol
×
g
mol
a. C2H4(g) + O3(g) → CH3CHO(g) + O2(g) ΔH° = −166 kJ − [143 kJ + 52 kJ] = −361 kJ
b. O3(g) + NO(g) → NO2(g) + O2(g) ΔH° = 34 kJ − [90. kJ + 143 kJ] = −199 kJ
c. SO3(g) + H2O(l) → H2SO4(aq) ΔH° = −909 kJ − [−396 kJ + (−286 kJ)] = −227 kJ
d. 2 NO(g) + O2(g) → 2 NO2(g)
ΔH° = 2(34) kJ − 2(90.) kJ = −112 kJ
ChemWork Problems
The answers to the problems 116-123 (or a variation to these problems) are found in OWL. These
problems are also assignable in OWL.
Challenge Problems
124.
Only when there is a volume change can PV work be done. In pathway 1 (steps 1 + 2), only
the first step does PV work (step 2 has a constant volume of 30.0 L). In pathway 2 (steps 3 +
4), only step 4 does PV work (step 3 has a constant volume of 10.0 L).
208
CHAPTER 6
THERMOCHEMISTRY
Pathway 1: w = −PΔV = −2.00 atm(30.0 L − 10.0 L) = −40.0 L atm ×
101.3 J
L atm
= −4.05 × 103 J
Pathway 2: w = −PΔV = −1.00 atm(30.0 L − 10.0 L) = −20.0 L atm ×
101.3 J
L atm
= −2.03 × 103 J
Note: The sign is negative because the system is doing work on the surroundings (an
expansion). We get different values of work for the two pathways; both pathways have the
same initial and final states. Because w depends on the pathway, work cannot be a state
function.
125.
A(l) → A(g)
ΔHvap = 30.7 kJ; at constant pressure, ΔH = qp = 30.7 kJ
Because PV = nRT, at constant pressure and temperature: w = −PΔV = −RTΔn, where:
Δn = moles of gaseous products − moles of gaseous reactants = 1 − 0 = 1
w = −RTΔn = −8.3145 J/K•mol(80. + 273 K)(1 mol) = −2940 J = −2.94 kJ
ΔE = q + w = 30.7 kJ + (−2.94 kJ) = 27.8 kJ
126.
Energy needed =
20. × 103 g C12 H 22 O11
1 mol C12 H 22 O11
5640 kJ
×
×
= 3.3 × 105 kJ/h
h
342.3 g C12 H 22 O11
mol
Energy from sun = 1.0 kW/m2 = 1000 W/m2 =
10,000 m2 ×
1.0 kJ
sm
2
×
Percent efficiency =
127.
1000 J
1.0 kJ
=
2
sm
s m2
60 s
60 min
= 3.6 × 107 kJ/h
×
min
h
3.3 × 105 kJ
energy used per hour
× 100 =
× 100 = 0.92%
total energy per hour
3.6 × 10 7 kJ
Energy used in 8.0 hours = 40. kWh =
Energy from the sun in 8.0 hours =
40. kJ h 3600 s
= 1.4 × 105 kJ
×
s
h
10. kJ 60 s 60 min
× 8.0 h = 2.9 × 104 kJ/m2
×
×
min
h
s m2
Only 19% of the sunlight is converted into electricity:
0.19 × (2.9 × 104 kJ/m2) × area = 1.4 × 105 kJ, area = 25 m2
128.
a. 2 HNO3(aq) + Na2CO3(s) → 2 NaNO3(aq) + H2O(l) + CO2(g)
ΔH° = [2(−467 kJ) + (−286 kJ) + (−393.5 kJ)] − [2(−207 kJ) + (−1131 kJ)] = −69 kJ
CHAPTER
6
THERMOCHEMISTRY
2.0 × 104 gallons ×
4 qt
946 mL
1.42 g
= 1.1 × 108 g of concentrated nitric
×
×
gal
qt
mL
acid solution
1.1 × 108 g solution ×
7.7 × 107 g HNO3 ×
7.7 × 107 g HNO3 ×
209
70.0 g HNO 3
= 7.7 × 107 g HNO3
100.0 g solution
1 mol HNO 3
63.02 g HNO 3
×
105.99 g Na 2 CO 3
1 mol Na 2 CO 3
×
mol Na 2 CO 3
2 mol HNO 3
= 6.5 × 107 g Na2CO3
1 mol HNO 3
− 69 kJ
= −4.2 × 107 kJ
×
63.02 g HNO 3
2 mol HNO 3
4.2 × 107 kJ of heat was released.
b. They feared the heat generated by the neutralization reaction would vaporize the
unreacted nitric acid, causing widespread airborne contamination.
129.
400 kcal ×
4.18 kJ
= 1.7 × 103 kJ ≈ 2 × 103 kJ
kcal


1 kg  9.81 m
2.54 cm
1m 
 = 160 J ≈ 200 J
 ×
PE = mgz = 180 lb ×
×  8 in ×
×
2
in
100 cm 
2.205 lb 
s


200 J of energy is needed to climb one step. The total number of steps to climb are:
2 × 106 J ×
130.
1 step
= 1 × 104 steps
200 J
H2(g) + 1/2 O2(g) → H2O(l) ΔH° = ΔH of , H 2O ( l ) = −285.8 kJ; we want the reverse reaction:
H2O(l) → H2(g) + 1/2 O2(g) ΔH° = 285.8 kJ
w = −P∆V; because PV = nRT, at constant T and P, P∆V = RT∆n, where ∆n = moles of
gaseous products – moles of gaseous reactants. Here, Δn = (1 mol H2 + 0.5 mol O2) – (0) =
1.50 mol.
ΔE° = ΔH° − PΔV = ΔH° − RTΔn


1 kJ
ΔE° = 285.8 kJ −  8.3145 J/K • mol × 298 K ×
× 1.50 mol 
1000 J


ΔE° = 285.8 kJ − 3.72 kJ = 282.1 kJ
131.
There are five parts to this problem. We need to calculate:
(1) q required to heat H2O(s) from −30. °C to 0°C; use the specific heat capacity of H2O(s)
(2) q required to convert 1 mol H2O(s) at 0°C into 1 mol H2O(l) at 0°C; use ∆Hfusion
210
CHAPTER 6
THERMOCHEMISTRY
(3) q required to heat H2O(l) from 0°C to 100.°C; use the specific heat capacity of H2O(l)
(4) q required to convert 1 mol H2O(l) at 100.°C into 1 mol H2O(g) at 100.°C;
use ∆Hvaporization
(5) q required to heat H2O(g) from 100.°C to 140.°C; use the specific heat capacity of
H2O(g)
We will sum up the heat required for all five parts, and this will be the total amount of heat
required to convert 1.00 mol of H2O(s) at −30.°C to H2O(g) at 140.°C.
q1 = 2.03 J/°C•g × 18.02 g × [0 – (−30.)]°C = 1.1 × 103 J
q2 = 1.00 mol × 6.02 × 103 J/mol = 6.02 × 103 J
q3 = 4.18 J/°C•g × 18.02 g × (100. – 0)°C = 7.53 × 103 J
q4 = 1.00 mol × 40.7 × 103 J/mol = 4.07 × 104 J
q5 = 2.02 J/°C•g × 18.02 g × (140. – 100.)°C = 1.5 × 103 J
qtotal = q1 + q2 + q3 + q4 + q5 = 5.69 × 104 J = 56.9 kJ
132.
When a mixture of ice and water exists, the temperature of the mixture remains at 0°C until
all of the ice has melted. Because an ice-water mixture exists at the end of the process, the
temperature remains at 0°C. All of the energy released by the element goes to convert ice into
water. The energy required to do this is related to ∆Hfusion = 6.02 kJ/mol (from Exercise 131).
Heat loss by element = heat gain by ice cubes at 0°C
Heat gain = 109.5 g H2O ×
1 mol H 2 O
6.02 kJ
= 36.6 kJ
×
mol H 2 O
18.02 g
Specific heat of element =
36,600 J
q
= 0.375 J/°C•g
=
mass × ΔT
500.0 g × (195 − 0) o C
Integrative Problems
133.
N2(g) + 2 O2(g) → 2 NO2(g)
∆H = 67.7 kJ
n N2 =
PV
3.50 atm × 0.250 L
= 2.86 × 10 −2 mol N2
=
0
.
08206
L atm
RT
× 373 K
K mol
n O2 =
PV
3.50 atm × 0.450 L
= 5.15 × 10 −2 mol O2
=
0.08206 L atm
RT
× 373 K
K mol
CHAPTER
6
THERMOCHEMISTRY
211
2.86 × 10 −2 mol N2 ×
2 mol NO 2
= 5.72 × 10 −2 mol NO2 produced if N2 is limiting.
1 mol N 2
5.15 × 10 −2 mol O2 ×
2 mol NO 2
= 5.15 × 10 −2 mol NO2 produced if O2 is limiting.
2 mol O 2
O2 is limiting because it produces the smaller quantity of product. The heat required is:
5.15 × 10 −2 mol NO2 ×
134.
67.7 kJ
= 1.74 kJ
2 mol NO 2
a. 4 CH3NO2(l) + 3 O2(g) → 4 CO2(g) + 2 N2(g) + 6 H2O(g)
ΔH orxn = −1288.5 kJ = [4 mol(−393.5 kJ/mol) + 6 mol(−242 kJ/mol)] −
[4 mol (∆H of , CH 3 NO 2 )]
Solving: ΔH of , CH 3 NO 2 = −434 kJ/mol
b.
Ptotal = 950. torr ×
n N2 =
0.168 atm × 15.0 L
= 0.0823 mol N2
0.08206 L atm
× 373 K
K mol
0.0823 mol N2 ×
135.
1 atm
= 1.25 atm; PN 2 = Ptotal × χ N 2 = 1.25 atm × 0.134
760 torr
= 0.168 atm
28.02 g N 2
= 2.31 g N2
1 mol N 2
Heat loss by U = heat gain by heavy water; volume of cube = (cube edge)3
Mass of heavy water = 1.00 × 103 mL ×
Heat gain by heavy water =
4.211 J
o
Heat loss by U = 1.4 × 104 J =
7.0 × 102 g U ×
1.11 g
= 1110 g
mL
× 1110 g × (28.5 – 25.5)°C = 1.4 × 104 J
Cg
0.117 J
o
× mass × (200.0 – 28.5)°C, mass = 7.0 × 102 g U
Cg
1 cm 3
= 37 cm3; cube edge = (37 cm3)1/3 = 3.3 cm
19.05 g
212
CHAPTER 6
THERMOCHEMISTRY
Marathon Problems
136.
Pathway I:
Step 1: (5.00 mol, 3.00 atm, 15.0 L) → (5.00 mol, 3.00 atm, 55.0 L)
w = −PΔV = −(3.00 atm)(55.0 L − 15.0 L) = −120. L atm
w = −120. L atm ×
1 kJ
101.3 J
×
= −12.2 kJ
L atm 1000 J
Step 1 is at constant pressure. The heat released/gained at constant pressure = qp = ΔH.
From the problem, ΔH = nCpΔT for an ideal gas. Using the ideal gas law, let’s substitute
for ΔT.
Δ(PV) = Δ(nRT) = nRΔT for a specific gas sample. So: ΔT =
Δ(PV)
nR
C Δ(PV)
Δ(PV)
; Note: Δ(PV) = (P2V2 − P1V1)
= p
nR
R
5
For an ideal monatomic gas, Cp = R; substituting into the above equation:
2
ΔH = qp = nCpΔT = nCp ×
5
 5  Δ(PV)
ΔH =  R 
= Δ (PV )
2
2  R
ΔH = qp =
5
5
Δ(PV) = (3.00 atm × 55.0 L − 3.00 atm × 15.0 L) = 300. L atm
2
2
ΔH = qp = 300. L atm ×
101.3 J
1 kJ
×
= 30.4 kJ
L atm
1000 J
ΔE = q + w = 30.4 kJ − 12.2 kJ = 18.2 kJ
Note: We could have used ΔE = nCvΔT to calculate the same answer (ΔE = 18.2 kJ).
Step 2: (5.00 mol, 3.00 atm, 55.0 L) → (5.00 mol, 6.00 atm, 20.0 L)
In this step, neither pressure nor volume are constant. We will need to determine q in a
different way. However, it will always hold for an ideal gas that ΔE = nCvΔT and ΔH =
nCpΔT.
 3   Δ(PV )  3
ΔE = nCvΔT = n  R  
 = ΔPV
 2   nR  2
3
ΔE = (120. − 165) L atm = −67.5 L atm (Carry an extra significant figure.)
2
ΔE = −67.5 L atm ×
101.3 J
1 kJ
×
= −6.8 kJ
L atm 1000 J
CHAPTER
6
THERMOCHEMISTRY
213
 5   Δ(PV )  5
ΔH = nCpΔT = n  R  
 = ΔPV
 2   nR  2
ΔH =
5
(120. − 165) L atm = −113 L atm (Carry an extra significant figure.)
2
w = −PΔV = −(6.00 atm)(20.0 − 55.0) L = 210. L atm
w = 210. L atm ×
1 kJ
101.3 J
×
= 21.3 kJ
L atm 1000 J
ΔE = q + w, −6.8 kJ = q + 21.3 kJ, q = −28.1 kJ
Summary:
Path I
q
w
ΔE
ΔH
Step 1
Step 2
30.4 kJ
−12.2 kJ
18.2 kJ
30.4 kJ
−28.1 kJ
21.3 kJ
−6.8 kJ
−11 kJ
Total
2.3 kJ
9.1 kJ
11.4 kJ
19 kJ
Pathway II:
Step 3: (5.00 mol, 3.00 atm, 15.0 L) → (5.00 mol, 6.00 atm, 15.0 L)
Step 3 is at constant volume. The heat released/gained at constant volume = qv = ΔE.
 3   Δ(PV )  3
ΔE = nCvΔT = n  R  
 = ΔPV
 2   nR  2
ΔE = qv =
3
3
Δ(PV) = (6.00 atm × 15.0 L − 3.00 atm × 15.0 L)
2
2
ΔE = qv =
3
(90.0 − 45.0) L atm = 67.5 L atm
2
ΔE = qv = 67.5 L atm ×
101.3 J
1 kJ
×
= 6.84 kJ
L atm 1000 J
w = −PΔV = 0 because ΔV = 0
ΔH = ΔE + Δ(PV) = 67.5 L atm + 45.0 L atm = 112.5 L atm = 11.40 kJ
Step 4: (5.00 mol, 6.00 atm, 15.0 L) → (5.00 mol, 6.00 atm, 20.0 L)
Step 4 is at constant pressure so qp = ΔH.
 5   Δ(PV )  5
ΔH = qp = nCpΔT =  R  
 = ΔPV
 2   nR  2
214
CHAPTER 6
ΔH =
THERMOCHEMISTRY
5
(120. − 90.0) L atm = 75 L atm
2
ΔH = qp = 75 L atm ×
101.3 J
1 kJ
×
= 7.6 kJ
L atm 1000 J
w = −PΔV = − (6.00 atm)(20.0 − 15.0) L = −30. L atm
w = −30. L atm ×
101.3 J
1 kJ
×
= −3.0 kJ
L atm 1000 J
ΔE = q + w = 7.6 kJ − 3.0 kJ = 4.6 kJ
Summary:
Path II
q
w
ΔE
ΔH
Step 3
6.84 kJ
0
6.84 kJ
11.40 kJ
Step 4
7.6 kJ
−3.0 kJ
4.6 kJ
7.6 kJ
Total
14.4 kJ
−3.0 kJ
11.4 kJ
19.0 kJ
State functions are independent of the particular pathway taken between two states; path
functions are dependent on the particular pathway. In this problem, the overall values of ΔH
and ΔE for the two pathways are the same (see the two summaries of results); hence ΔH and
ΔE are state functions. However, the overall values of q and w for the two pathways are
different; hence q and w are path functions.
137.
 2 x + y/2 
CxHy + 
 Ο2 → x CO2 + y/2 H2O
2


[x(−393.5) + y/2 (−242)] − ΔH oC x H y = −2044.5, − (393.5)x − 121y − ΔH C x H y = −2044.5
dgas =
P • MM
, where MM = average molar mass of CO2/H2O mixture
RT
0.751 g/L =
1.00 atm × MM
, MM of CO2/H2O mixture = 29.1 g/mol
0.08206 L atm
× 473 K
K mol
Let a = mol CO2 and 1.00 − a = mol H2O (assuming 1.00 total moles of mixture)
(44.01)a + (1.00 − a) × 18.02 = 29.1; solving: a = 0.426 mol CO2 ; mol H2O = 0.574 mol
y
0.574
y
Thus:
= 2 , 2.69 = , y = (2.69)x
0.426
x
x
For whole numbers, multiply by three, which gives y = 8, x = 3. Note that y = 16, x = 6 is
possible, along with other combinations. Because the hydrocarbon has a lower density than
Kr, the molar mass of CxHy must be less than the molar mass of Kr (83.80 g/mol). Only C3H8
works.
CHAPTER
6
THERMOCHEMISTRY
−2044.5 = −393.5(3) − 121(8) − ΔH oC3H8 , ΔH oC3H8 = −104 kJ/mol
215
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
Questions
19.
The equations relating the terms are νλ = c, E = hν, and E = hc/λ. From the equations,
wavelength and frequency are inversely related, photon energy and frequency are directly
related, and photon energy and wavelength are inversely related. The unit of 1 Joule (J) = 1
kg m2/s2. This is why you must change mass units to kg when using the deBroglie equation.
20.
Frequency is the number of waves (cycles) of electromagnetic radiation per second that pass
a given point in space. Speed refers to the distance a wave travels per unit time. All
electromagnetic radiation (EMR) travels at the same speed (c, the speed of light = 2.998 × 108
m/s). However, each wavelength of EMR has its own unique frequency,
21.
The photoelectric effect refers to the phenomenon in which electrons are emitted from the
surface of a metal when light strikes it. The light must have a certain minimum frequency
(energy) in order to remove electrons from the surface of a metal. Light having a frequency
below the minimum value results in no electrons being emitted, whereas light at or higher
than the minimum frequency does cause electrons to be emitted. For light having a frequency
higher than the minimum frequency, the excess energy is transferred into kinetic energy for
the emitted electron. Albert Einstein explained the photoelectric effect by applying quantum
theory.
22.
The emission of light by excited atoms has been the key interconnection between the
macroscopic world we can observe and measure, and what is happening on a microscopic
basis within an atom. Excited atoms emit light (which we can observe and measure) because
of changes in the microscopic structure of the atom. By studying the emissions of atoms, we
can trace back to what happened inside the atom. Specifically, our current model of the atom
relates the energy of light emitted to electrons in the atom moving from higher allowed
energy states to lower allowed energy states.
23.
Example 7.3 calculates the deBroglie wavelength of a ball and of an electron. The ball has a
wavelength on the order of 10 −34 m. This is incredibly short and, as far as the wave- particle
duality is concerned, the wave properties of large objects are insignificant. The electron, with
its tiny mass, also has a short wavelength; on the order of 10 −10 m. However, this wavelength
is significant because it is on the same order as the spacing between atoms in a typical crystal.
For very tiny objects like electrons, the wave properties are important. The wave properties
must be considered, along with the particle properties, when hypothesizing about the electron
motion in an atom.
215
216
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
24. a. For hydrogen (Z = 1), the energy levels in units of joules are given by the equation En =
−2.178 × 10−18(1/n2). As n increases, the differences between 1/n2 for consecutive energy
levels becomes smaller and smaller. Consider the difference between 1/n2 values for n = 1
and n = 2 as compared to n = 3 and n = 4.
For n = 1 and n = 2:
1
2
1
−
1
2
2
For n = 3 and n = 4:
= 1 − 0.25 = 0.75
1
3
2
−
1
42
= 0.1111 – 0.0625 = 0.0486
Because the energy differences between 1/n2 values for consecutive energy levels decrease as
n increases, the energy levels get closer together as n increases.
b. For a spectral transition for hydrogen, ΔE = Ef − Ei:
 1
1 
ΔE = −2.178 × 10 −18 J  2 − 2 
 nf
ni 

where ni and nf are the levels of the initial and final states, respectively. A positive value of
ΔE always corresponds to an absorption of light, and a negative value of ΔE always
corresponds to an emission of light.
In the diagram, the red line is for the ni = 3 to nf = 2 transition.
1 
 1
ΔE = −2.178 × 10 −18 J  2 − 2  = −2.178 × 10 −18
3 
2
1 1
J − 
4 9
ΔE = −2.178 × 10 −18 J × (0.2500 − 0.1111) = −3.025 × 10 −19 J
The photon of light must have precisely this energy (3.025 × 10 −19 J).
|ΔE| = Ephoton = hν =
λ =
hc
λ
hc
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
=
= 6.567 × 10 −7 m = 656.7 nm
−19
| ΔE |
3.025 × 10 J
From Figure 7.2, λ = 656.7 nm is red light so the diagram is correct for the red line.
In the diagram, the green line is for the ni = 4 to nf = 2 transition.
1 
 1
ΔE = −2.178 × 10 −18 J  2 − 2  = −4.084 × 10 −19 J
4 
2
λ =
hc
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
=
= 4.864 × 10 −7 m = 486.4 nm
−
19
| ΔE |
4.084 × 10 J
From Figure 7.2, λ = 486.4 nm is green-blue light. The diagram is consistent with this
line.
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
217
In the diagram, the blue line is for the ni = 5 to nf = 2 transition.
1 
 1
ΔE = −2.178 × 10 −18 J  2 − 2  = −4.574 × 10 −19 J
5 
2
λ =
hc
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
=
= 4.343 × 10 −7 m = 434.3 nm
| ΔE |
4.574 × 10 −19 J
From Figure 7.2, λ = 434.3 nm is blue or blue-violet light. The diagram is consistent
with this line also.
25.
The Bohr model was an important step in the development of the current quantum
mechanical model of the atom. The idea that electrons can only occupy certain, allowed
energy levels is illustrated nicely (and relatively easily). We talk about the Bohr model to
present the idea of quantized energy levels.
26.
The figure on the left tells us that the probability of finding the electron in the 1s orbital at
points along a line drawn outward from the nucleus in any direction. This probability is
greatest close to the nucleus and drops off rapidly as the distance from the nucleus increases.
The figure on the right represents the total probability of finding the electron at a particular
distance from the nucleus for a 1s hydrogen orbital. For this distribution, the hydrogen 1s
orbital is divided into successive thin spherical shells and the total probability of finding the
electron in each spherical shell is plotted versus distance from the nucleus. This graph is
called the radial probability distribution.
The radial probability distribution initially shows a steady increase with distance from the
nucleus, reaches a maximum, then shows a steady decrease. Even though it is likely to find an
electron near the nucleus, the volume of the spherical shell close to the nucleus is tiny,
resulting in a low radial probability. The maximum radial probability distribution occurs at a
distance of 5.29 × 10 −2 nm from the nucleus; the electron is most likely to be found in the
volume of the shell centered at this distance from the nucleus. The 5.29 × 10 −2 nm distance is
the exact radius of innermost (n = 1) orbit in the Bohr model.
27.
When the p and d orbital functions are evaluated at various points in space, the results
sometimes have positive values and sometimes have negative values. The term phase is often
associated with the + and − signs. For example, a sine wave has alternating positive and
negative phases. This is analogous to the positive and negative values (phases) in the p and d
orbitals.
28.
The widths of the various blocks in the periodic table are determined by the number of
electrons that can occupy the specific orbital(s). In the s block, we have one orbital (ℓ = 0, mℓ
= 0) that can hold two electrons; the s block is two elements wide. For the f block, there are 7
degenerate f orbitals (ℓ = 3, mℓ = −3, −2, −1, 0, 1, 2, 3), so the f block is 14 elements wide.
The g block corresponds to ℓ = 4. The number of degenerate g orbitals is 9. This comes from
the 9 possible mℓ values when ℓ = 4 (mℓ = −4, −3, −2, −1, 0, 1, 2, 3, 4). With 9 orbitals, each
orbital holding two electrons, the g block would be 18 elements wide. The h block has ℓ = 5,
mℓ = −5, −4, −3, −2, −1, 0, 1, 2, 3, 4, 5. With 11 degenerate h orbitals, the h block would be
22 elements wide.
218
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
29.
If one more electron is added to a half-filled subshell, electron-electron repulsions will
increase because two electrons must now occupy the same atomic orbital. This may slightly
decrease the stability of the atom.
30.
Size decreases from left to right and increases going down the periodic table. Thus, going one
element right and one element down would result in a similar size for the two elements
diagonal to each other. The ionization energies will be similar for the diagonal elements since
the periodic trends also oppose each other. Electron affinities are harder to predict, but atoms
with similar sizes and ionization energies should also have similar electron affinities.
31.
The valence electrons are strongly attracted to the nucleus for elements with large ionization
energies. One would expect these species to readily accept another electron and have very
exothermic electron affinities. The noble gases are an exception; they have a large ionization
energy but have an endothermic electron affinity. Noble gases have a filled valence shell of
electrons. The added electron in a noble gas must go into a higher n value atomic orbital,
having a significantly higher energy, and this is very unfavorable.
32.
Electron-electron repulsions become more important when we try to add electrons to an atom.
From the standpoint of electron-electron repulsions, larger atoms would have more favorable
(more exothermic) electron affinities. Considering only electron-nucleus attractions, smaller
atoms would be expected to have the more favorable (more exothermic) electron affinities.
These trends are exactly the opposite of each other. Thus the overall variation in electron
affinity is not as great as ionization energy in which attractions to the nucleus dominate.
33.
For hydrogen and hydrogen-like (one-electron ions), all atomic orbitals with the same n value
have the same energy. For polyatomic atoms/ions, the energy of the atomic orbitals also
depends on ℓ. Because there are more nondegenerate energy levels for polyatomic atoms/ions
as compared to hydrogen, there are many more possible electronic transitions resulting in
more complicated line spectra.
34.
Each element has a characteristic spectrum because each element has unique energy levels.
Thus the presence of the characteristic spectral lines of an element confirms its presence in
any particular sample.
35.
Yes, the maximum number of unpaired electrons in any configuration corresponds to a
minimum in electron-electron repulsions.
36.
The electron is no longer part of that atom. The proton and electron are completely separated.
37.
Ionization energy applies to the removal of the electron from an atom in the gas phase. The
work function applies to the removal of an electron from the surface of a solid element.
M(g) → M+(g) + e− ionization energy; M(s) → M+(s) + e− work function
38.
Li+ ions are the smallest of the alkali metal cations and will be most strongly attracted to the
water molecules.
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
219
Exercises
Light and Matter
c
ν =
40.
99.5 MHz = 99.5 × 106 Hz = 99.5 × 106 s −1 ; λ =
41.
ν=
λ
c
λ
=
=
2.998 × 108 m/s
39.
= 3.84 × 1014 s−1
780. × 10 −9 m
3.00 × 108 m/s
−2
1.0 × 10 m
c
2.998 × 108 m / s
= 3.01 m
=
v
99.5 × 10 6 s −1
= 3.0 × 1010 s −1
E = hν = 6.63 × 10-34 J s × 3.0 × 1010 s −1 = 2.0 × 10 −23 J/photon
2.0 × 10 −23 J 6.02 × 10 23 photons
×
= 12 J/mol
photon
mol
42.
E = hν =
hc
λ
=
6.63 × 10 −34 J s × 3.00 × 108 m/s
= 8.0 × 10 −18 J/photon
1m
25 nm ×
1 × 109 nm
6.02 × 10 23 photons
8.0 × 10 −18 J
×
= 4.8 × 106 J/mol
mol
photon
43.
280 nm: ν =
320 nm: ν =
c
λ
=
3.00 × 108 m/s
= 1.1 × 1015 s −1
1m
280 nm ×
1 × 109 nm
3.00 × 108 m/s
= 9.4 × 1014 s −1
320 × 10 −9 nm
The compounds in the sunscreen absorb ultraviolet B (UVB) electromagnetic radiation
having a frequency from 9.4 × 1014 s −1 to 1.1 × 1015 s −1 .
44.
S-type cone receptors: λ =
λ =
c
ν
=
2.998 × 108 m/s
6.00 × 10 s
14
2.998 × 108 m/s
7.49 × 10 s
14
−1
−1
= 5.00 × 10 −7 m = 500. nm
= 4.00 × 10 −7 m = 400. nm
S-type cone receptors detect 400-500 nm light. From Figure 7.2 in the text, this is violet
to green light, respectively.
220
CHAPTER 7
M-type cone receptors: λ =
λ =
ATOMIC STRUCTURE AND PERIODICITY
2.998 × 108 m/s
4.76 × 1014 s −1
2.998 × 108 m/s
6.62 × 10 s
14
−1
= 6.30 × 10 −7 m = 630. nm
= 4.53 × 10 −7 m = 453 nm
M-type cone receptors detect 450-630 nm light. From Figure 7.2 in the text, this is blue
to orange light, respectively.
L-type cone receptors: λ =
λ =
2.998 × 108 m/s
4.28 × 10 s
14
−1
2.998 × 108 m/s
6.00 × 10 s
14
−1
= 7.00 × 10 −7 m = 700. nm
= 5.00 × 10 −7 m = 500. nm
L-type cone receptors detect 500-700 nm light.
respectively.
45.
This represents green to red light,
The wavelength is the distance between consecutive wave peaks. Wave a shows 4 wavelengths, and wave b shows 8 wavelengths.
Wave a: λ =
1.6 × 10 −3 m
= 4.0 × 10−4 m
4
Wave b: λ =
1.6 × 10 −3 m
= 2.0 × 10−4 m
8
Wave a has the longer wavelength. Because frequency and photon energy are both inversely
proportional to wavelength, wave b will have the higher frequency and larger photon energy
since it has the shorter wavelength.
ν =
E =
c
λ
hc
λ
=
=
2.998 × 108 m/s
2.0 × 10
−4
= 1.5 × 1012 s−1
m
6.626 × 10 −34 J s × 2.998 × 108 m/s
2.0 × 10
−4
= 9.9 × 10−22 J
m
Because both waves are examples of electromagnetic radiation, both waves travel at the same
speed, c, the speed of light. From Figure 7.2 of the text, both of these waves represent
infrared electromagnetic radiation.
46.
Referencing Figure 7.2 of the text, 2.12 × 10−10 m electromagnetic radiation is X rays.
λ=
c
2.9979 × 108 m / s
=
= 2.799 m
ν
107.1 × 10 6 s −1
From the wavelength calculated above, 107.1 MHz electromagnetic radiation is FM radiowaves.
CHAPTER 7
λ =
ATOMIC STRUCTURE AND PERIODICITY
221
hc
6.626 × 10 −34 J s × 2.998 × 108 m / s
=
= 5.00 × 10−7 m
E
3.97 × 10 −19 J
The 3.97 × 10−19 J/photon electromagnetic radiation is visible (green) light.
The photon energy and frequency order will be the exact opposite of the wavelength ordering
because E and ν are both inversely related to λ. From the previously calculated wavelengths,
the order of photon energy and frequency is:
FM radiowaves < visible (green) light < X rays
longest λ
shortest λ
lowest ν
highest ν
smallest E
largest E
47.
Ephoton =
hc
λ
=
1.98 × 105 J ×
48.
Ephoton = hν =
6.626 × 10 −34 J s × 2.998 × 108 m/s
= 1.32 × 10 −18 J
1m
150. nm ×
1 × 109 nm
1 photon
1.32 × 10
hc
λ
−18
J
, E photon =
×
1 atom C
= 1.50 × 1023 atoms C
photon
6.626 × 10 −34 J s × 2.998 × 108 m/s
1.0 × 10
−10
= 2.0 × 10−15 J
m
2.0 × 10 −15 J 6.02 × 10 23 photons
1 kJ
×
×
= 1.2 × 106 kJ/mol
photon
mol
1000 J
Ephoton =
6.626 × 10 −34 J s × 2.998 × 108 m/s
1.0 × 10 m
4
= 2.0 × 10−29 J
2.0 × 10 −29 J 6.02 × 10 23 photons
1 kJ
×
×
= 1.2 × 10−8 kJ/mol
photon
mol
1000 J
X rays do have an energy greater than the carbon-carbon bond energy. Therefore, X rays
could conceivably break carbon-carbon bonds in organic compounds and thereby disrupt the
function of an organic molecule. Radiowaves, however, do not have sufficient energy to
break carbon-carbon bonds and are therefore relatively harmless.
49.
The energy needed to remove a single electron is:
279.7 kJ
1 mol
= 4.645 × 10 −22 kJ = 4.645 × 10 −19 J
×
mol
6.0221 × 10 23
E=
hc
λ
, λ=
hc
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
= 4.277 × 10 −7 m = 427.7 nm
=
−19
E
4.645 × 10 J
222
50.
CHAPTER 7
208.4 kJ
1 mol
= 3.461 × 10 −22 kJ = 3.461 × 10 −19 J to remove one electron
×
23
mol
6.0221 × 10
E=
51.
hc
λ
, λ=
hc
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
= 5.739 × 10 −7 m = 573.9 nm
=
E
3.461 × 10 −19 J
Ionization energy = energy to remove an electron = 7.21 × 10−19 = Ephoton
Ephoton = hν and λν = c. So ν =
λ=
52.
ATOMIC STRUCTURE AND PERIODICITY
hc
E photon
=
c
λ
and E =
hc
λ
.
6.626 × 10 −34 J s × 2.998 × 108 m/s
7.21 × 10
−19
= 2.76 × 10−7 m = 276 nm
J
1.478 × 10 −21 kJ 1.478 × 10 −18 J
890.1 kJ
1 mol
×
=
=
atom
atom
mol
6.0221 × 10 23 atoms
= ionization energy per atom
E=
hc
λ
,λ =
hc
6.626 × 10 −34 J s × 2.9979 × 108 m/s
= 1.344 × 10−7 m = 134.4 nm
=
−18
E
1.478 × 10 J
No, it will take light having a wavelength of 134.4 nm or less to ionize gold. A photon of
light having a wavelength of 225 nm is longer wavelength and thus lower energy than 134.4
nm light.
53.
a.
10.% of speed of light = 0.10 × 3.00 × 108 m/s = 3.0 × 107 m/s
λ=
h
6.63 × 10 −34 J s
= 2.4 × 10 −11 m = 2.4 × 10 −2 nm
, λ =
mv
9.11 × 10 −31 kg × 3.0 × 10 7 m/s
Note: For units to come out, the mass must be in kg because 1 J =
b.
λ=
1 kg m 2
s2
.
h
6.63 × 10 −34 J s
=
= 3.4 × 10 −34 m = 3.4 × 10 −25 nm
mv
0.055 kg × 35 m/s
This number is so small that it is insignificant. We cannot detect a wavelength this small.
The meaning of this number is that we do not have to worry about the wave properties of
large objects.
54.
a.
λ=
h
6.626 × 10 −34 J s
=
= 1.32 × 10 −13 m
mv 1.675 × 10 − 27 kg × (0.0100 × 2.998 × 108 m/s)
b.
λ=
h
h
6.626 × 10 −34 J s
, v=
=
= 5.3 × 103 m/s
−12
− 27
mv
λm
75 × 10 m × 1.675 × 10 kg
CHAPTER 7
55.
λ=
ATOMIC STRUCTURE AND PERIODICITY
223
h
h
6.63 × 10 −34 J s
=
= 1.6 × 10 −27 kg
, m=
mv
λv 1.5 × 10 −15 m × (0.90 × 3.00 × 108 m/s)
This particle is probably a proton or a neutron.
56.
λ=
h
h
; for λ = 1.0 × 102 nm = 1.0 × 10 −7 m:
, v=
mv
λm
6.63 × 10 −34 J s
= 7.3 × 103 m/s
−31
−7
9.11 × 10 kg × 1.0 × 10 m
v =
For λ = 1.0 nm = 1.0 × 10 −9 m: v =
6.63 × 10 −34 J s
= 7.3 × 105 m/s
9.11 × 10 −31 kg × 1.0 × 10 −9 m
Hydrogen Atom: The Bohr Model
57.
For the H atom (Z = 1): En = −2.178 × 10-18 J/n 2; for a spectral transition, ΔE = Ef − Ei:
 1
1 
ΔE = −2.178 × 10 −18 J  2 − 2 
n
ni 
 f
where ni and nf are the levels of the initial and final states, respectively. A positive value of
ΔE always corresponds to an absorption of light, and a negative value of ΔE always
corresponds to an emission of light.
1 
 1
a. ΔE = −2.178 × 10 −18 J  2 − 2  = −2.178 × 10 −18
3 
2
1
1
J − 
9
4

ΔE = −2.178 × 10 −18 J × (0.2500 − 0.1111) = −3.025 × 10 −19 J
The photon of light must have precisely this energy (3.025 × 10 −19 J).
|ΔE| = Ephoton = hν =
hc
λ
, λ =
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
hc
=
| ΔE |
3.025 × 10 −19 J
= 6.567 × 10 −7 m = 656.7 nm
From Figure 7.2, this is visible electromagnetic radiation (red light).
1 
 1
b. ΔE = −2.178 × 10 −18 J  2 − 2  = −4.084 × 10 −19 J
4 
2
λ =
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
hc
=
= 4.864 × 10 −7 m = 486.4 nm
| ΔE |
4.084 × 10 −19 J
This is visible electromagnetic radiation (green-blue light).
224
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
1 
1
c. ΔE = −2.178 × 10 −18 J  2 − 2  = −1.634 × 10 −18 J
2 
1
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
λ=
1.634 × 10
−18
J
= 1.216 × 10 −7 m = 121.6 nm
This is ultraviolet electromagnetic radiation.
58.
1 
 1
a. ΔE = −2.178 × 10 −18 J  2 − 2  = −1.059 × 10 −19 J
4 
3
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
hc
=
= 1.876 × 10 −6 m = 1876 nm
| ∆E |
1.059 × 10 −19 J
λ=
From Figure 7.2, this is infrared electromagnetic radiation.
1 
 1
b. ΔE = −2.178 × 10 −18 J  2 − 2  = −4.901 × 10 −20 J
5 
4
λ =
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
hc
=
= 4.053 × 10 −6 m
− 20
| ΔE |
4.901 × 10
J
= 4053 nm (infrared)
1 
 1
c. ΔE = −2.178 × 10 −18 J  2 − 2  = −1.549 × 10 −19 J
5 
3
λ =
59.
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
hc
=
= 1.282 × 10 −6 m
−
19
| ΔE |
1.549 × 10 J
= 1282 nm (infrared)
5
4
3
a. 3 → 2
a
b. 4 → 2
b
E 2
c. 2 → 1
c
Energy levels are not to scale.
n=1
60.
5
4
3
E 2
n=1
a. 4 → 3
b
a
c
b. 5 → 4
c. 5 → 3
Energy levels are not to scale.
CHAPTER 7
61.
ATOMIC STRUCTURE AND PERIODICITY
225
The longest wavelength light emitted will correspond to the transition with the smallest
energy change (smallest ∆E). This is the transition from n = 6 to n = 5.
1 
 1
ΔE = −2.178 × 10−18 J  2 − 2  = −2.662 × 10−20 J
6 
5
λ =
hc
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
= 7.462 × 10−6 m = 7462 nm
=
− 20
| ∆E |
2.662 × 10
J
The shortest wavelength emitted will correspond to the largest ΔE; this is n = 6 → n = 1.
1 
1
ΔE = −2.178 × 10−18 J  2 − 2  = −2.118 × 10−18 J
6 
1
λ =
62.
63.
hc
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
= 9.379 × 10−8 m = 93.79 nm
=
−
18
| ∆E |
2.118 × 10 J
There are 4 possible transitions for an electron in the n = 5 level (5 → 4, 5 → 3, 5 → 2, and 5
→ 1). If an electron initially drops to the n = 4 level, three additional transitions can occur (4
→ 3, 4 → 2, and 4 → 1). Similarly, there are two more transitions from the n = 3 level (3 →
2, 3 → 1) and one more transition for the n = 2 level (2 → 1). There are a total of 10 possible
transitions for an electron in the n = 5 level for a possible total of 10 different wavelength
emissions.
 1
1 
ΔE = −2.178 × 10−18 J  2 − 2  = −2.178 × 10−18 J
ni 
 nf
λ =
 1
1

− 2  = 2.091 × 10−18 J = Ephoton
2
5
1 

hc
6.6261 × 10 −34 J s × 2.9979 × 108 m / s
=
= 9.500 × 10−8 m = 95.00 nm
E
2.091 × 10 −18 J
Because wavelength and energy are inversely related, visible light (λ ≈ 400−700 nm) is not
energetic enough to excite an electron in hydrogen from n = 1 to n = 5.
 1
1 
ΔE = −2.178 × 10−18 J  2 − 2  = 4.840 × 10−19 J
6
2 

λ =
hc
6.6261 × 10 −34 J s × 2.9979 × 108 m / s
=
= 4.104 × 10−7 m = 410.4 nm
−19
E
4.840 × 10 J
Visible light with λ = 410.4 nm will excite an electron from the n = 2 to the n = 6 energy
level.
64.
a. False; it takes less energy to ionize an electron from n = 3 than from the ground state.
b. True
226
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
c. False; the energy difference between n = 3 and n = 2 is smaller than the energy
difference between n = 3 and n = 1; thus the wavelength is larger for the n = 3 → n = 2
electronic transition than for the n = 3 → n = 1 transition. E and λ are inversely
proportional to each other (E = hc/λ).
d. True
e. False; n = 2 is the first excited state, and n = 3 is the second excited state.
65.
Ionization from n = 1 corresponds to the transition ni = 1 → nf = ∞, where E∞ = 0.
1
ΔE = E∞ − E1 = −E1 = 2.178 × 10 −18  2  = 2.178 × 10 −18 J = Ephoton
1 
λ =
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
hc
=
= 9.120 × 10 −8 m = 91.20 nm
E
2.178 × 10 −18 J
 1 
To ionize from n = 2, ΔE = E∞ − E2 = −E2 = 2.178 × 10 −18  2  = 5.445 × 10 −19 J
2 
λ =
66.
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
5.445 × 10
−19
= 3.648 × 10 −7 m = 364.8 nm
J
 1 
ΔE = E∞ − En = −En = 2.178 × 10 −18 J  2 
n 
Ephoton =
6.626 × 10 −34 J s × 2.9979 × 108 m/s
hc
=
= 1.36 × 10 −19 J
−9
λ
1460 × 10 m
 1 
Ephoton = ΔE = 1.36 × 10 −19 J = 2.178 × 10 −18  2 , n2 = 16.0, n = 4
n 
67.
|ΔE| = Ephoton = hν = 6.662 × 10 −34 J s × 6.90 × 1014 s −1 = 4.57 × 10 −19 J
ΔE = −4.57 × 10 −19 J because we have an emission.
1 
 1
−4.57 × 10 −19 J = En – E5 = −2.178 × 10 −18 J  2 − 2 
5 
n
1
n
2
−
1
= 0.210,
25
1
= 0.250, n2 = 4, n = 2
n2
The electronic transition is from n = 5 to n = 2.
68.
|ΔE| = Ephoton =
hc
λ
=
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
397.2 × 10
−9
m
ΔE = −5.001 × 10 −19 J because we have an emission.
= 5.001 × 10 −19 J
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
227
1 
 1
−5.001 × 10 −19 J = E2 − En = −2.178 × 10 −18 J  2 − 2 
n 
2
0.2296 =
1
1
1
− 2 , 2 = 0.0204, n = 7
4
n
n
Quantum Mechanics, Quantum Numbers, and Orbitals
69.
a. Δp = mΔv = 9.11 × 10−31 kg × 0.100 m/s =
ΔpΔx ≥
b.
Δx =
9.11 × 10 −32 kg m
s
h
h
6.626 × 10 −34 J s
=
, Δx =
= 5.79 × 10−4 m
4π
4 π Δp
4 × 3.142 × (9.11 × 10 −32 kg m / s)
h
6.626 × 10 −34 J s
=
= 3.64 × 10−33 m
4 × 3.142 × 0.145 kg × 0.100 m / s
4 π Δp
c. The diameter of an H atom is roughly ∼10−8 cm. The uncertainty in position is much
larger than the size of the atom.
d. The uncertainty is insignificant compared to the size of a baseball.
70.
Units of ΔE • Δt = J × s, the same as the units of Planck's constant.
Units of Δ(mv) • Δx = kg ×
m
kg m 2
kg m 2
× m=
=
×s = J×s
s
s
s2
71.
n = 1, 2, 3, ... ; ℓ = 0, 1, 2, ... (n − 1); mℓ = −ℓ ... -2, -1, 0, 1, 2, ...+ℓ
72.
a. This general shape represents a p orbital (ℓ = 1) and because there is a node in each of the
lobes, this figure represents a 3p orbital (n = 3, ℓ = 1)
b. This is an s orbital (ℓ = 0). And because there is one node present, this is a 2s orbital
(n = 2, ℓ = 0).
c. This is the shape of a specific d oriented orbital (ℓ = 2). This orbital is designated as a
d z 2 . Because no additional nodes are present inside any of the boundary surfaces, this is
a 3d z 2 orbital (n = 3, ℓ = 2).
73.
a. allowed
b. For ℓ = 3, mℓ can range from −3 to +3; thus +4 is not allowed.
c. n cannot equal zero.
74.
a. For n = 3, ℓ = 3 is not possible.
d. ms cannot equal −1.
e. ℓ cannot be a negative number.
d. ℓ cannot be a negative number.
228
CHAPTER 7
f.
ATOMIC STRUCTURE AND PERIODICITY
For ℓ = 1, mℓ cannot equal 2.
The quantum numbers in parts b and c are allowed.
75.
ψ2 gives the probability of finding the electron at that point.
76.
The diagrams of the orbitals in the text give only 90% probabilities of where the electron may
reside. We can never be 100% certain of the location of the electrons due to Heisenberg’s
uncertainty principle.
Polyelectronic Atoms
77.
He: 1s2; Ne: 1s22s22p6; Ar: 1s22s22p63s23p6; each peak in the diagram corresponds to a
subshell with different values of n. Corresponding subshells are closer to the nucleus for
heavier elements because of the increased nuclear charge.
78.
In polyelectronic atoms, the orbitals of a given principal quantum level are not degenerate. In
polyelectronic atoms, the energy order of the n = 1, 2, and 3 orbitals are (not to scale):
3d
3p
E
3s
2p
2s
1s
In general, the lower the n value for an orbital, the closer on average the electron can be to the
nucleus, and the lower the energy. Within a specific n value orbital (like 2s vs. 2p or 3s vs.
3p vs. 3d), it is generally true that Ens < Enp < End < Enf.
To rationalize this order, we utilize the radial probability distributions. In the 2s and 2p
distribution, notice that the 2s orbital has a small hump of electron density very near the
nucleus. This indicates that an electron in the 2s orbital can be very close to the nucleus some
of the time. The 2s electron penetrates to the nucleus more than a 2p electron, and with this
penetration comes a lower overall energy for the 2s orbital as compared to the 2p orbital.
In the n = 3 radial probability distribution, the 3s electron has two humps of electron density
very close to the nucleus, and the 3p orbital has one hump very close to the nucleus. The 3s
orbital electron is most penetrating, with the 3p orbital electron the next most penetrating,
followed by the least penetrating 3d orbital electron. The more penetrating the electron, the
lower the overall energy. Hence the 3s orbital is lower energy than the 3p orbitals which is
lower energy than the 3d orbitals.
CHAPTER 7
79.
ATOMIC STRUCTURE AND PERIODICITY
3d z 2 : one orbital
5p: three orbitals
229
4d: five orbitals
n = 5: ℓ = 0 (1 orbital), ℓ = 1 (3 orbitals), ℓ = 2 (5 orbitals), ℓ = 3 (7 orbitals),
ℓ = 4 (9 orbitals); total for n = 5 is 25 orbitals.
n = 4: ℓ = 0 (1), ℓ = 1 (3), ℓ = 2 (5), ℓ = 3 (7); total for n = 4 is 16 orbitals.
80.
1p, 0 electrons (ℓ ≠ 1 when n = 1); 6d x 2 − y 2 , 2 electrons (specifies one atomic orbital); 4f, 14
electrons (7 orbitals have 4f designation); 7py, 2 electrons (specifies one atomic orbital); 2s,
2 electrons (specifies one atomic orbital); n = 3, 18 electrons (3s, 3p, and 3d orbitals are
possible; there are one 3s orbital, three 3p orbitals, and five 3d orbitals).
81.
a. n = 4: ℓ can be 0, 1, 2, or 3. Thus we have s (2 e− ), p (6 e− ), d (10 e− ), and f (14 e− )
orbitals present. Total number of electrons to fill these orbitals is 32.
b. n = 5, mℓ = +1: For n = 5, ℓ = 0, 1, 2, 3, 4. For ℓ = 1, 2, 3, 4, all can have mℓ = +1. There
are four distinct orbitals having these quantum numbers, which can hold 8 electrons.
c. n = 5, ms = +1/2: For n = 5, ℓ = 0, 1, 2, 3, 4. Number of orbitals = 1, 3, 5, 7, 9 for each
value of ℓ, respectively. There are 25 orbitals with n = 5. They can hold 50 electrons, and
25 of these electrons can have ms = +1/2.
d. n = 3, ℓ = 2: These quantum numbers define a set of 3d orbitals. There are 5 degenerate
3d orbitals that can hold a total of 10 electrons.
e. n = 2, ℓ = 1: These define a set of 2p orbitals. There are 3 degenerate 2p orbitals that can
hold a total of 6 electrons.
82.
a. It is impossible to have n = 0. Thus no electrons can have this set of quantum numbers.
b. The four quantum numbers completely specify a single electron in a 2p orbital.
c. n = 3, ms = +1/2: 3s, 3p, and 3d orbitals all have n = 3. These nine orbitals can each hold
one electron with ms = +1/2; 9 electrons can have these quantum numbers
d. n = 2, ℓ = 2: this combination is not possible (ℓ ≠ 2 for n = 2). Zero electrons in an atom
can have these quantum numbers.
e. n = 1, ℓ = 0, mℓ = 0: these define a 1s orbital that can hold 2 electrons.
83.
a. Na: 1s22s22p63s1; Na has 1 unpaired electron.
or
1s
2s
2p
3s
3s
230
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
b. Co: 1s22s22p63s23p64s23d7; Co has 3 unpaired electrons.
1s
2s
2p
3s
3p
or
4s
3d
3d
c. Kr: 1s22s22p63s23p64s23d104p6; Kr has 0 unpaired electrons.
1s
2s
4s
84.
2p
3s
3d
3p
4p
The two exceptions are Cr and Cu.
Cr: 1s22s22p63s23p64s13p5; Cr has 6 unpaired electrons.
1s
2s
2p
3p
or
or
4s
3s
4s
3d
3d
Cu: 1s22s22p63s23p64s13d10; Cu has 1 unpaired electron.
1s
2s
2p
3s
3p
or
4s
85.
4s
3d
Si: 1s22s22p63s23p2 or [Ne]3s23p2; Ga: 1s22s22p63s23p64s23d104p1 or [Ar]4s23d104p1
As: [Ar]4s23d104p3; Ge: [Ar]4s23d104p2; Al: [Ne]3s23p1; Cd: [Kr]5s24d10
S: [Ne]3s23p4; Se: [Ar]4s23d104p4
CHAPTER 7
86.
ATOMIC STRUCTURE AND PERIODICITY
231
Cu: [Ar]4s23d9 (using periodic table), [Ar]4s13d10 (actual)
O: 1s22s22p4; La: [Xe]6s25d1; Y: [Kr]5s24d1; Ba: [Xe]6s2
Tl: [Xe]6s24f145d106p1; Bi: [Xe]6s24f145d106p3
87.
The following are complete electron configurations. Noble gas shorthand notation could also
be used.
Sc: 1s22s22p63s23p64s23d1; Fe: 1s22s22p63s23p64s23d6
P:
1s22s22p63s23p3; Cs: 1s22s22p63s23p64s23d104p65s24d105p66s1
Eu: 1s22s22p63s23p64s23d104p65s24d105p66s24f65d1*
Pt: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d8*
Xe: 1s22s22p63s23p64s23d104p65s24d105p6; Br: 1s22s22p63s23p64s23d104p5
*Note: These electron configurations were predicted using only the periodic table.
The actual electron configurations are: Eu: [Xe]6s24f7 and Pt: [Xe]6s14f145d9
88.
Cl: ls22s22p63s23p5 or [Ne]3s23p5
Sb: [Kr]5s24d105p3
Sr: 1s22s22p63s23p64s23d104p65s2 or [Kr]5s2
W: [Xe]6s24f145d4
Pb: [Xe]6s24f145d106p2
Cf: [Rn]7s25f10*
*Note: Predicting electron configurations for lanthanide and actinide elements is difficult
since they have 0, 1, or 2 electrons in d orbitals. This is the actual Cf electron configuration.
89.
a. Both In and I have one unpaired 5p electron, but only the nonmetal I would be expected
to form a covalent compound with the nonmetal F. One would predict an ionic
compound to form between the metal In and the nonmetal F.
I: [Kr]5s24d105p5
↑↓ ↑↓ ↑
5p
b. From the periodic table, this will be element 120. Element 120: [Rn]7s25f146d107p68s2
c. Rn: [Xe]6s24f145d106p6; note that the next discovered noble gas will also have 4f electrons
(as well as 5f electrons).
d. This is chromium, which is an exception to the predicted filling order. Cr has 6 unpaired
electrons, and the next most is 5 unpaired electrons for Mn.
Cr: [Ar]4s13d5 ↑ ↑ ↑ ↑ ↑ ↑
4s
3d
232
90.
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
a. As: 1s22s22p63s23p64s23d104p3
b. Element 116 will be below Po in the periodic table: [Rn]7s25f146d107p4
c. Ta: [Xe]6s24f145d3 or Ir: [Xe]6s24f145d7
d. At: [Xe]6s24f145d106p5; note that element 117 will also have electrons in the 6p atomic
orbitals (as well as electrons in the 7p orbitals).
91.
a. The complete ground state electron for this neutral atom is 1s22s22p63s23p4. This
atom has 2 + 2 + 6 + 2 + 4 = 16 electrons. Because the atom is neutral, it also has 16
protons, making the atom sulfur, S.
b. Complete excited state electron configuration: 1s22s12p4; this neutral atom has 2 + 1 + 4
= 7 electrons, which means it has 7 protons, which identifies it as nitrogen, N.
c. Complete ground state electron configuration: 1s22s22p63s23p64s23d104p5; this 1−
charged ion has 35 electrons. Because the overall charge is 1−, this ion has 34 protons
which identifies it as selenium. The ion is Se−.
92.
a. This atom has 10 electrons. Ne
b. S
c. The predicted ground state configuration is [Kr]5s24d9. From the periodic table, the
element is Ag. Note: [Kr]5s14d10 is the actual ground state electron configuration for Ag.
d. Bi: [Xe]6s24f145d106p3; the three unpaired electrons are in the 6p orbitals.
93.
Hg: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d10
a. From the electron configuration for Hg, we have 3s2, 3p6, and 3d10 electrons; 18 total
electrons with n = 3.
b. 3d10, 4d10, 5d10; 30 electrons are in the d atomic orbitals.
c. 2p6, 3p6, 4p6, 5p6; each set of np orbitals contain one pz atomic orbital. Because we have
4 sets of np orbitals and two electrons can occupy the pz orbital, there are 4(2) = 8
electrons in pz atomic orbitals.
d. All the electrons are paired in Hg, so one-half of the electrons are spin up (ms = +1/2) and
the other half are spin down (ms = −1/2). 40 electrons have spin up.
94.
Element 115, Uup, is in Group 5A under Bi (bismuth):
Uup: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d107p3
a. 5s2, 5p6, 5d10, and 5f14; 32 electrons have n = 5 as one of their quantum numbers.
b. ℓ = 3 are f orbitals. 4f14 and 5f14 are the f orbitals used. They are all filled, so 28 electrons
have ℓ = 3.
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
233
c. p, d, and f orbitals all have one of the degenerate orbitals with mℓ = 1. There are 6 orbitals
with mℓ = 1 for the various p orbitals used; there are 4 orbitals with mℓ =1 for the various
d orbitals used; and there are 2 orbitals with mℓ = 1 for the various f orbitals used. We
have a total of 6 + 4 + 2 = 12 orbitals with mℓ = 1. Eleven of these orbitals are filled with
2 electrons, and the 7p orbitals are only half-filled. The number of electrons with mℓ = 1
is 11 × (2 e−) + 1 × (1 e−) = 23 electrons.
d. The first 112 electrons are all paired; one-half of these electrons (56 e−) will have ms =
−1/2. The 3 electrons in the 7p orbitals singly occupy each of the three degenerate 7p
orbitals; the three electrons are spin parallel, so the 7p electrons either have ms = +1/2 or
ms = −1/2. Therefore, either 56 electrons have ms = −1/2 or 59 electrons have ms = −1/2.
95.
B: 1s22s22p1
1s
1s
2s
2s
2p*
n
ℓ
mℓ
1
1
2
2
2
0
0
0
0
1
0
0
0
0
−1
ms
+1/2
−1/2
+1/2
−1/2
+1/2
*This is only one of several possibilities for the 2p electron. The 2p electron in B could have
mℓ = −1, 0 or +1 and ms = +1/2 or −1/2 for a total of six possibilities.
N: 1s22s22p3
1s
1s
2s
2s
2p
2p
2p
96.
n
ℓ
1
1
2
2
2
2
2
0
0
0
0
1
1
1
Ti : [Ar]4s23d2
n ℓ mℓ
mℓ
0
0
0
0
−1
0
+1
ms
+1/2
−1/2
+1/2
−1/2
+1/2
+1/2
+1/2
(Or all 2p electrons could have ms = −1/2.)
ms
4s
4s
3d
4
4
3
0 0 +1/2
0 0 −1/2
2 −2 +1/2
3d
3
2
−1 +1/2
Only one of 10 possible combinations of mℓ and ms for the first d
electron. For the ground state, the second d electron should be in
a different orbital with spin parallel; 4 possibilities.
234
97.
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
Group 1A: 1 valence electron; ns1; Li: [He]2s1; 2s1 is the valence electron configuration for
Li.
Group 2A: 2 valence electrons; ns2; Ra: [Rn]7s2; 7s2 is the valence electron configuration
for Ra.
Group 3A: 3 valence electrons; ns2np1; Ga: [Ar]4s23d104p1; 4s24p1 is the valence electron
configuration for Ga. Note that valence electrons for the representative elements of Groups
1A-8A are considered those electrons in the highest n value, which for Ga is n = 4. We do
not include the 3d electrons as valence electrons because they are not in n = 4 level.
Group 4A: 4 valence electrons; ns2np2; Si: [Ne]3s23p2; 3s23p2 is the valence electron
configuration for Si.
Group 5A: 5 valence electrons; ns2np3; Sb: [Kr]5s24d105p3; 5s25p3 is the valence electron
configuration for Sb.
Group 6A: 6 valence electrons; ns2np4; Po: [Xe]6s24f145d106p4; 6s26p4 is the valence
electron configuration for Po.
Group 7A: 7 valence electrons; ns2np5; 117: [Rn]7s25f146d107p5; 7s27p5 is the valence
electron configuration for 117.
Group 8A: 8 valence electrons; ns2np6; Ne: [He]2s22p6; 2s22p6 is the valence electron
configuration for Ne.
98.
a.
2 valence electrons; 4s2
b. 6 valence electrons; 2s22p4
c.
7 valence electrons; 7s27p5
d. 3 valence electrons; 5s25p1
e.
8 valence electrons; 3s23p6
f.
5 valence electrons; 6s26p3
99.
O: 1s22s22px22py2 (↑↓ ↑↓ __ ); there are no unpaired electrons in this oxygen atom. This
configuration would be an excited state, and in going to the more stable ground state
(↑↓ ↑ ↑ ), energy would be released.
100.
The number of unpaired electrons is in parentheses.
a. excited state of boron
B ground state: 1s22s22p1
c. exited state of fluorine
F ground state: 1s22s22p5
↑↓ ↑↓ ↑
2p
(1)
(1)
(3)
(1)
b. ground state of neon
(0)
Ne ground state: 1s22s22p6 (0)
d. excited state of iron
(6)
Fe ground state: [Ar]4s23d6 (4)
↑↓ ↑
↑ ↑
3d
↑
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
235
101.
None of the s block elements have 2 unpaired electrons. In the p block, the elements with
either ns2np2 or ns2np4 valence electron configurations have 2 unpaired electrons. For
elements 1-36, these are elements C, Si, and Ge (with ns2np2) and elements O, S, and Se (with
ns2np4). For the d block, the elements with configurations nd2 or nd8 have two unpaired
electrons. For elements 1-36, these are Ti (3d2) and Ni (3d8). A total of 8 elements from the
first 36 elements have two unpaired electrons in the ground state.
102.
The s block elements with ns1 for a valence electron configuration have one unpaired
electron. These are elements H, Li, Na, and K for the first 36 elements. The p block elements
with ns2np1 or ns2np5 valence electron configurations have one unpaired electron. These are
elements B, Al, and Ga (ns2np1) and elements F, Cl, and Br (ns2np5) for the first 36 elements.
In the d block, Sc ([Ar]4s23d1) and Cu ([Ar]4s13d10) each have one unpaired electron. A total
of 12 elements from the first 36 elements have one unpaired electron in the ground state.
103.
We get the number of unpaired electrons by examining the incompletely filled subshells. The
paramagnetic substances have unpaired electrons, and the ones with no unpaired electrons are
not paramagnetic (they are called diamagnetic).
Li: 1s22s1 ↑ ; paramagnetic with 1 unpaired electron.
2s
N: 1s22s22p3 ↑ ↑ ↑ ; paramagnetic with 3 unpaired electrons.
2p
Ni: [Ar]4s23d8 ↑↓ ↑↓ ↑↓ ↑ ↑ ; paramagnetic with 2 unpaired electrons.
3d
Te: [Kr]5s24d105p4 ↑↓ ↑ ↑ ; paramagnetic with 2 unpaired electrons.
5p
Ba: [Xe]6s2 ↑↓ ; not paramagnetic because no unpaired electrons are present.
6s
Hg: [Xe]6s24f145d10 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ; not paramagnetic because no unpaired electrons.
5d
104.
We get the number of unpaired electrons by examining the incompletely filled subshells.
O: [He]2s22p4
2p4:
↑↓ ↑
O+: [He]2s22p3
2p3:
↑
O−: [He]2s22p5
2p5:
↑↓ ↑↓ ↑
Os: [Xe]6s24f145d6
5d6:
↑↓ ↑
Zr: [Kr]5s24d2
4d2:
↑
S: [Ne]3s23p4
3p4:
↑↓ ↑
two unpaired e−
↑
three unpaired e−
↑ ↑
↑
one unpaired e−
↑ ↑
four unpaired e−
two unpaired e−
↑
↑
two unpaired e−
236
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
F: [He]2s22p5
2p5:
↑↓ ↑↓ ↑
one unpaired e−
Ar: [Ne]3s23p6
3p6
↑↓ ↑↓ ↑↓
zero unpaired e−
The Periodic Table and Periodic Properties
105.
Size (radius) decreases left to right across the periodic table, and size increases from top to
bottom of the periodic table.
a. S < Se < Te
b. Br < Ni < K
c. F < Si < Ba
All follow the general radius trend.
106.
a. Be < Na < Rb
b. Ne < Se < Sr
c. O < P < Fe
All follow the general radius trend.
107.
The ionization energy trend is the opposite of the radius trend; ionization energy (IE), in
general, increases left to right across the periodic table and decreases from top to bottom of
the periodic table.
a. Te < Se < S
b. K < Ni < Br
c. Ba < Si < F
All follow the general ionization energy trend.
108.
a. Rb < Na < Be
b. Sr < Se < Ne
c. Fe < P < O
All follow the general ionization energy trend.
109.
a. He (From the general radius trend.)
b. Cl
c. Element 116 is the next oxygen family member to be discovered (under Po), element 119
is the next alkali metal to be discovered (under Fr), and element 120 is the next alkaline
earth metal to be discovered (under Ra). From the general radius trend, element 116 will
be the smallest.
d. Si
e. Na+; this ion has the fewest electrons as compared to the other sodium species present.
Na+ has the smallest number of electron-electron repulsions, which makes it the smallest
ion with the largest ionization energy.
110.
a.
Ba (From the general ionization energy trend.)
b.
K
c. O; in general, Group 6A elements have a lower ionization energy than neighboring
Group 5A elements. This is an exception to the general ionization energy trend across
the periodic table.
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
237
d. S2−; this ion has the most electrons compared to the other sulfur species present. S2− has
the largest number of electron-electron repulsions, which leads to S2− having the largest
size and smallest ionization energy.
e. Cs; this follows the general ionization energy trend.
111.
a. Sg: [Rn]7s25f146d4
b. W
c. Sg is in the same group as chromium and would be expected to form compounds and ions
similar to that of chromium. CrO3, Cr2O3, CrO42−, Cr2O72− are some know chromium
compounds/ions, so SgO3, Sg2O3, SgO42−, and Sg2O72− are some likely possibilities.
112.
a. Uus will have 117 electrons. [Rn]7s25f146d107p5
b. It will be in the halogen family and will be most similar to astatine (At).
c. Uus should form 1− charged anions like the other halogens. Like the other halogens,
some possibilities are NaUus, Mg(Uus)2, C(Uus)4, and O(Uus)2
d. Assuming Uus is like the other halogens, some possibilities are UusO-, UusO2-, UusO3-,
and UusO4-.
113.
As: [Ar]4s23d104p3; Se: [Ar]4s23d104p4; the general ionization energy trend predicts that Se
should have a higher ionization energy than As. Se is an exception to the general ionization
energy trend. There are extra electron-electron repulsions in Se because two electrons are in
the same 4p orbital, resulting in a lower ionization energy for Se than predicted.
114.
Expected order from the ionization energy trend: Be < B < C < N < O
B and O are exceptions to the general ionization energy trend. The ionization energy of O is
lower because of the extra electron-electron repulsions present when two electrons are paired
in the same orbital. This makes it slightly easier to remove an electron from O compared to
N. B is an exception because of the smaller penetrating ability of the 2p electron in B
compared to the 2s electrons in Be. The smaller penetrating ability makes it slightly easier to
remove an electron from B compared to Be. The correct ionization energy ordering, taking
into account the two exceptions, is B < Be < C < O < N.
115.
a. As we remove succeeding electrons, the electron being removed is closer to the nucleus,
and there are fewer electrons left repelling it. The remaining electrons are more strongly
attracted to the nucleus, and it takes more energy to remove these electrons.
b. Al: 1s22s22p63s23p1; for I4, we begin removing an electron with n = 2. For I3, we remove
an electron with n = 3 (the last valence electron). In going from n = 3 to n = 2, there is a
big jump in ionization energy because the n = 2 electrons are closer to the nucleus on
average than the n = 3 electrons. Since the n = 2 electrons are closer, on average, to the
nucleus, they are held more tightly and require a much larger amount of energy to remove
compared to the n = 3 electrons. In general, valence electrons are much easier to remove
than inner-core electrons.
238
116.
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
The general ionization energy trend says that ionization energy increases going left to right
across the periodic table. However, one of the exceptions to this trend occurs between
Groups 2A and 3A. Between these two groups, Group 3A elements usually have a lower
ionization energy than Group 2A elements. Therefore, Al should have the lowest first
ionization energy value, followed by Mg, with Si having the largest ionization energy.
Looking at the values for the first ionization energy in the graph, the green plot is Al, the blue
plot is Mg, and the red plot is Si.
Mg (the blue plot) is the element with the huge jump between I2 and I3. Mg has two valence
electrons, so the third electron removed is an inner core electron. Inner core electrons are
always much more difficult to remove than valence electrons since they are closer to the
nucleus, on average, than the valence electrons.
117.
a. More favorable electron affinity: C and Br; the electron affinity trend is very erratic.
Both N and Ar have positive electron affinity values (unfavorable) due to their electron
configurations (see text for detailed explanation).
b. Higher ionization energy: N and Ar (follows the ionization energy trend)
c. Larger size: C and Br (follows the radius trend)
118.
a. More favorable electron affinity: K and Cl; Mg has a positive electron affinity value,
and F has a more positive electron affinity value than expected from its position relative
to Cl.
b. Higher ionization energy: Mg and F
119.
c. Larger radius: K and Cl
Al(−44), Si(−120), P(−74), S(−200.4), Cl(−348.7); based on the increasing nuclear charge,
we would expect the electron affinity values to become more exothermic as we go from left
to right in the period. Phosphorus is out of line. The reaction for the electron affinity of P is:
P(g) + e− → P−(g)
[Ne]3s23p3
[Ne]3s23p4
The additional electron in P− will have to go into an orbital that already has one electron.
There will be greater repulsions between the paired electrons in P−, causing the electron
affinity of P to be less favorable than predicted based solely on attractions to the nucleus.
120.
Electron affinity refers to the energy associated with the process of adding an electron to a
gaseous substance. Be, N, and Ne all have endothermic (unfavorable) electron affinity values.
In order to add an electron to Be, N, or Ne, energy must be added. Another way of saying
this is that Be, N, and Ne become less stable (have a higher energy) when an electron is
added to each. To rationalize why those three atoms have endothermic (unfavorable) electron
affinity values, let’s see what happens to the electron configuration as an electron is added.
Be(g) +
[He]2s2
e− → Be−(g)
[He]2s22p1
Ne(g) + e− → Ne−(g)
[He]2s22p6
[He]2s22p63s1
N(g) + e− → N−(g)
[He]2s22p3
[He]2s22p4
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
239
In each case something energetically unfavorable occurs when an electron is added. For Be,
the added electron must go into a higher-energy 2p atomic orbital because the 2s orbital is
full. In N, the added electron must pair up with another electron in one of the 2p atomic
orbitals; this adds electron-electron repulsions. In Ne, the added electron must be added to a
much higher 3s atomic orbital because the n = 2 orbitals are full.
121.
The electron affinity trend is very erratic. In general, electron affinity decreases down the
periodic table, and the trend across the table is too erratic to be of much use.
a. Se < S; S is most exothermic.
b. I < Br < F < Cl; Cl is most exothermic.
(F is an exception).
122.
a. N < O < F, F is most exothermic.
b. Al < P < Si; Si is most exothermic.
123.
Electron-electron repulsions are much greater in O− than in S− because the electron goes into
a smaller 2p orbital versus the larger 3p orbital in sulfur. This results in a more favorable
(more exothermic) electron affinity for sulfur.
124.
O; the electron-electron repulsions will be much more severe for O− + e− → O2− than for O
+ e− → O−, resulting in O having the more exothermic (favorable) electron affinity.
125.
a. Se3+(g) → Se4+(g) + e−
b. S−(g) + e− → S2−(g)
c. Fe3+(g) + e− → Fe2+(g)
d. Mg(g) → Mg+(g) + e−
126.
a. The electron affinity (EA) of Mg2+ is ΔH for Mg2+(g) + e− → Mg+(g); this is just the
reverse of the second ionization energy (I2) for Mg. EA(Mg2+) = −I2(Mg) = −1445
kJ/mol (Table 7.5). Note that when an equation is reversed, the sign on the equation is
also reversed.
b. I1 of Cl− is ΔH for Cl−(g) → Cl(g) + e−; I1(Cl− ) = −EA(Cl) = 348.7 kJ/mol (Table 7.7)
c. Cl+(g) + e− → Cl(g)
ΔH = − I1(Cl) = −1255 kJ/mol = EA(Cl+) (Table 7.5)
d. Mg−(g) → Mg(g) + e− ΔH = −EA(Mg) = −230 kJ/mol = I1(Mg− )
Alkali Metals
127.
It should be potassium peroxide (K2O2) because K+ ions are stable in ionic compounds. K2+
ions are not stable; the second ionization energy of K is very large compared to the first.
128.
a. Li3N; lithium nitride
129.
ν=
c
λ
=
b. NaBr; sodium bromide
2.9979 × 108 m/s
455.5 × 10
−9
m
= 6.582 × 1014 s −1
E = hν = 6.6261 × 10 −34 J s × 6.582 × 1014 s −1 = 4.361 × 10 −19 J
c. K2S; potassium sulfide
240
130.
CHAPTER 7
For 589.0 nm: ν =
c
λ
=
ATOMIC STRUCTURE AND PERIODICITY
2.9979 × 108 m/s
589.0 × 10 −9 m
= 5.090 × 1014 s −1
E = hν = 6.6261 × 10 −34 J s × 5.090 × 1014 s −1 = 3.373 × 10 −19 J
For 589.6 nm: ν = c/λ = 5.085 × 1014 s −1 ; E = hν = 3.369 × 10 −19 J
The energies in kJ/mol are:
3.373 × 10 −19 J ×
1 kJ
6.0221 × 10 23
×
= 203.1 kJ/mol
1000 J
mol
3.369 × 10 −19 J ×
1 kJ
6.0221 × 10 23
×
= 202.9 kJ/mol
1000 J
mol
131.
Yes; the ionization energy general trend is to decrease down a group, and the atomic radius
trend is to increase down a group. The data in Table 7.8 confirm both of these general trends.
132.
It should be element 119 with the ground state electron configuration [Rn]7s25f146d107p68s1.
133.
a. 6 Li(s) + N2(g) → 2 Li3N(s)
b. 2 Rb(s) + S(s) → Rb2S(s)
134.
a. 2 Cs(s) + 2 H2O(l) → 2 CsOH(aq) + H2(g)
b. 2 Na(s) + Cl2(g) → 2 NaCl(s)
Additional Exercises
135.
No; lithium metal is very reactive. It will react somewhat violently with water, making it
completely unsuitable for human consumption. Lithium has a small first ionization energy,
so it is more likely that the lithium prescribed will be in the form of a soluble lithium salt (a
soluble ionic compound with Li+ as the cation).
136.
a. λ =
c
ν
=
3.00 × 108 m/s
13 −1
6.0 × 10 s
= 5.0 × 10 −6 m
b. From Figure 7.2, this is infrared electromagnetic radiation.
c. E = hν = 6.63 × 10 −34 J s × 6.0 × 1013 s −1 = 4.0 × 10 −20 J/photon
4.0 × 10 −20 J
6.022 × 10 23 photons
= 2.4 × 104 J/mol
×
photon
mol
d. Frequency and photon energy are directly related (E = hv). Because 5.4 × 1013 s −1
electromagnetic radiation (EMR) has a lower frequency than 6.0 × 1013 s −1 EMR, the 5.4
× 1013 s −1 EMR will have less energetic photons.
CHAPTER 7
137.
138.
ATOMIC STRUCTURE AND PERIODICITY
241
E=
310 kJ
1 mol
= 5.15 × 10 −22 kJ = 5.15 × 10 −19 J
×
23
mol
6.022 × 10
E =
6.626 × 10 −34 J s × 2.998 × 108 m/s
hc
hc
=
= 3.86 × 10 −7 m = 386 nm
, λ=
λ
E
5.15 × 10 −19 J
Energy to make water boil = s × m × ΔT =
Ephoton =
hc
λ
=
4.18 J
× 50.0 g × 75.0°C = 1.57 × 104 J
°C g
6.626 × 10 −34 J s × 2.998 × 108 m/s
9.75 × 10
−2
= 2.04 × 10−24 J
m
1.57 × 104 J ×
1 photon
1s
= 20.9 s; 1.57 × 104 J ×
= 7.70 × 1027 photons
− 24
750. J
2.04 × 10
J
139.
60 × 106 km ×
1000 m
1s
= 200 s (about 3 minutes)
×
km
3.00 × 108 m
140.
λ=
100 cm
hc
6.626 × 10 −34 J s × 2.998 × 108 m/s
= 5.53 × 10 −7 m ×
=
−
19
m
E
3.59 × 10 J
= 5.53 × 10 −5 cm
From the spectrum, λ = 5.53 × 10 −5 cm is greenish-yellow light.
141.
 1
1 
1 
 1
ΔE = −RH  2 − 2  = −2.178 × 10 −18 J  2 − 2  = −4.840 × 10 −19 J
n
6 
ni 
2
 f
λ=
100 cm
hc
6.6261 × 10 −34 J s × 2.9979 × 108 m / s
= 4.104 × 10 −7 m ×
=
−
19
m
| ΔE |
4.840 × 10 J
= 4.104 × 10 −5 cm
From the spectrum, λ = 4.104 × 10 −5 cm is violet light, so the n = 6 to n = 2 visible spectrum
line is violet.
142.
Exceptions: Cr, Cu, Nb, Mo, Tc, Ru, Rh, Pd, Ag, Pt, and Au; Tc, Ru, Rh, Pd, and Pt do not
correspond to the supposed extra stability of half-filled and filled subshells as normally
applied.
143.
a. True for H only.
144.
n = 5; mℓ = −4, −3, −2, −1, 0, 1, 2, 3, 4; 18 electrons
145.
1p: n = 1, ℓ = 1 is not possible; 3f: n = 3, ℓ = 3 is not possible; 2d: n = 2, ℓ = 2 is not
possible; in all three incorrect cases, n = ℓ. The maximum value ℓ can have is n − 1, not n.
146.
O: 1s22s22p4; C: 1s22s22p2; H: 1s1; N: 1s22s22p3; Ca: [Ar]4s2; P: [Ne]3s22p3
b. True for all atoms.
c. True for all atoms.
242
147.
148.
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
Mg: [Ne]3s2; K: [Ar]4s1
From the radii trend, the smallest-size element (excluding hydrogen) would be the one in the
most upper right corner of the periodic table. This would be O. The largest-size element
would be the one in the most lower left of the periodic table. Thus K would be the largest.
The ionization energy trend is the exact opposite of the radii trend. So K, with the largest
size, would have the smallest ionization energy. From the general ionization energy trend, O
should have the largest ionization energy. However, there is an exception to the general
ionization energy trend between N and O. Due to this exception, N would have the largest
ionization energy of the elements examined.
a. The 4+ ion contains 20 electrons. Thus the electrically neutral atom will contain 24
electrons. The atomic number is 24, which identifies it as chromium.
b. The ground state electron configuration of the ion must be 1s22s22p63s23p64s03d2; there
are 6 electrons in s orbitals.
c. 12
d. 2
e. From the mass, this is the isotope
f.
50
24 Cr.
There are 26 neutrons in the nucleus.
1s22s22p63s23p64s13d5 is the ground state electron configuration for Cr.
exception to the normal filling order.
Cr is an
149.
Valence electrons are easier to remove than inner-core electrons. The large difference in
energy between I2 and I3 indicates that this element has two valence electrons. This element
is most likely an alkaline earth metal since alkaline earth metal elements all have two valence
electrons.
150.
All oxygen family elements have ns2np4 valence electron configurations, so this nonmetal is
from the oxygen family.
a. 2 + 4 = 6 valence electrons.
b. O, S, Se, and Te are the nonmetals from the oxygen family (Po is a metal).
c. Because oxygen family nonmetals form 2− charged ions in ionic compounds, K2X would
be the predicted formula, where X is the unknown nonmetal.
d. From the size trend, this element would have a smaller radius than barium.
e. From the ionization energy trend, this element would have a smaller ionization energy
than fluorine.
151.
a.
Na(g) → Na+(g) + e−
Cl(g) + e− → Cl−(g)
Na(g) + Cl(g) → Na+(g) + Cl−g)
I1 = 495 kJ
EA = −348.7 kJ
ΔH = 146 kJ
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
Mg(g) → Mg+(g) + e−
F(g) + e−→ F-(g)
b.
I1 = 735 kJ
EA = −327.8 kJ
Mg(g) + F(g) → Mg+(g) + F−(g)
ΔH = 407 kJ
Mg+(g) → Mg2+(g) + e−
F(g) + e − → F−(g)
c.
243
I2 = 1445 kJ
EA = −327.8 kJ
Mg+(g) + F(g) → Mg2+(g) + F−(g)
ΔH = 1117 kJ
d. Using parts b and c, we get:
152.
Mg(g) + F(g) → Mg+(g) + F−(g)
Mg+(g) + F(g) → Mg2+(g) + F−(g)
ΔH = 407 kJ
ΔH = 1117 kJ
Mg(g) + 2 F(g) → Mg2+(g) + 2 F−(g)
ΔH = 1524 kJ
Sc: [Ar]4s23d1, 1 unpaired electron; Ti: [Ar]4s23d2; 2 unpaired e−; V: [Ar]4s23d3, 3 unpaired
e−; Cr: [Ar]4s13d5, 6 unpaired electrons (Cr is an exception to the normal filling order);
Mn: [Ar]4s23d5, 5 unpaired e−; Fe: [Ar]4s23d6, 4 unpaired e−; Co: [Ar]4s23d7, 3 unpaired e−;
Ni: [Ar]4s23d8, 2 unpaired e−; Cu: [Ar] 4s13d10, 1 unpaired e− (Cu is also an exception to
the normal filling order); Zn: [Ar]4s23d10, 0 unpaired e−.
ChemWork Problems
The answers to the problems 153-162 (or a variation to these problems) are found in OWL. These
problems are also assignable in OWL.
Challenge Problems
163.
λ =
h
, where m = mass and v = velocity; v rms =
mv
For one atom, R =
2.31 × 10−11 m =
Molar mass =
3RT
, λ =
m
h
m
3RT
m
=
h
3RTm
1 mol
8.3145 J
= 1.381 × 10−23 J/K•atom
×
23
K mol
6.022 × 10 atoms
6.626 × 10 −34 J s
m 3(1.381 × 10
− 23
26
23
, m = 5.32 × 10− kg = 5.32 × 10− g
)(373 K )
5.32 × 10 −23 g 6.022 × 10 23 atoms
= 32.0 g/mol
×
atom
mol
The atom is sulfur (S).
244
164.
CHAPTER 7
Ephoton =
hc
λ
=
ATOMIC STRUCTURE AND PERIODICITY
6.6261 × 10 −34 J s × 2.9979 × 108 m/s
253.4 × 10 −9 m
= 7.839 × 10 −19 J
ΔE = 7.839 × 10 −19 J; the general energy equation for one-electron ions is En = −2.178 ×
10 −18 J (Z2)/n2, where Z = atomic number.
 1
1 
ΔE = −2.178 × 10 −18 J (Z)2  2 − 2  , Z = 4 for Be3+
n
ni 
 f
 1
1 
ΔE = −7.839 × 10 −19 J = −2.178 × 10 −18 (4)2  2 − 2 
nf
5 

7.839 × 10 −19
2.178 × 10
−18
× 16
+
1
1
1
= 2 , 2 = 0.06249, nf = 4
25
nf nf
This emission line corresponds to the n = 5 → n = 4 electronic transition.
165.
a. Because wavelength is inversely proportional to energy, the spectral line to the right of B
(at a larger wavelength) represents the lowest possible energy transition; this is n = 4 to n
= 3. The B line represents the next lowest energy transition, which is n = 5 to n = 3, and
the A line corresponds to the n = 6 to n = 3 electronic transition.
b. Because this spectrum is for a one-electron ion, En = −2.178 × 10−18 J (Z2/n2). To
determine ΔE and, in turn, the wavelength of spectral line A, we must determine Z, the
atomic number of the one electron species. Use spectral line B data to determine Z.
 Z2
Z2 
ΔE5 → 3 = −2.178 × 10−18 J  2 − 2  = −2.178 × 10−18
5 
3
E =
 16 Z 2 


 9 × 25 


hc
6.6261 × 10 −34 J s(2.9979 × 108 m / s)
=
= 1.394 × 10−18 J
−9
λ
142.5 × 10 m
Because an emission occurs, ΔE5 → 3 = −1.394 × 10−18 J.
 16 Z 2 
 , Z2 = 9.001, Z = 3; the ion is Li2+.
ΔE = −1.394 × 10−18 J = −2.178 × 10−18 J 

9
×
25


Solving for the wavelength of line A:
1 
 1
ΔE6 → 3 = −2.178 × 10−18(3)2  2 − 2  = −1.634 × 10−18 J
6 
3
6.6261 × 10 −34 J s(2.9979 × 108 m / s)
hc
=
= 1.216 × 10−7 m = 121.6 nm
λ =
−18
ΔE
1.634 × 10 J
CHAPTER 7
166.
ATOMIC STRUCTURE AND PERIODICITY
245
1 
 1
For hydrogen: ΔE = −2.178 × 10-18 J  2 − 2  = −4.574 × 10 −19 J
5 
2
For a similar blue light emission, He+ will need about the same ΔE value.
For He+: En = −2.178 × 10 −18 J (Z2/n2), where Z = 2:
 22
22 
ΔE = −4.574 × 10 −19 J = −2.178 × 10-18 J  2 − 2 
n
4 
 f
0.2100 =
4
nf2
−
4
4
, 0.4600 = 2 ,
16
nf
nf = 2.949
The transition from n = 4 to n = 3 for He+ should emit a similar colored blue light as the n
= 5 to n = 2 hydrogen transition; both these transitions correspond to very nearly the
same energy change.
167.
For one-electron species, En = − R H Z 2 /n 2 . The ground state ionization energy is the energy
change for the n = 1 → n = ∞ transition. So:
ionization energy = E∞ − E1 = −E1 = R H Z 2 /n 2 = RHZ2
4.72 × 10 4 kJ
1 mol
1000 J
×
×
= 2.178 × 10−18 J (Z2); solving: Z = 6
23
mol
kJ
6.022 × 10
Element 6 is carbon (X = carbon), and the charge for a one-electron carbon ion is
5+ (m = 5). The one-electron ion is C5+.
168.
A node occurs when ψ = 0. ψ300 = 0 when 27 − 18σ + 2σ2 = 0.
Solving using the quadratic formula: σ =
18 ± (18) 2 − 4(2)(27)
4
=
18 ± 108
4
σ = 7.10 or σ = 1.90; because σ = r/ao, the nodes occur at r = (7.10)ao = 3.76 × 10−10 m and at
r = (1.90)ao = 1.01 × 10−10 m, where r is the distance from the nucleus.
169.
For r = ao and θ = 0° (Z = 1 for H):
ψ 2pz =
1
4(2π )
1/2


1


11
−
 5.29 × 10 


3/2
(1) e −1/2 cos 0 = 1.57 × 1014; ψ2 = 2.46 × 1028
For r = ao and θ = 90°, ψ 2 p z = 0 since cos 90° = 0; ψ2 = 0; there is no probability of finding
an electron in the 2pz orbital with θ = 0°. As expected, the xy plane, which corresponds to θ
= 0°, is a node for the 2pz atomic orbital.
246
170.
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
a. Each orbital could hold 3 electrons.
b. The first period corresponds to n = 1 which can only have 1s orbitals. The 1s orbital
could hold 3 electrons; hence the first period would have three elements. The second
period corresponds to n = 2, which has 2s and 2p orbitals. These four orbitals can each
hold three electrons. A total of 12 elements would be in the second period.
c. 15
171 .
d. 21
a. 1st period:
p = 1, q = 1, r = 0, s = ±1/2 (2 elements)
2nd period:
p = 2, q = 1, r = 0, s = ±1/2 (2 elements)
3rd period:
p = 3, q = 1, r = 0, s = ±1/2 (2 elements)
p = 3, q = 3, r = −2, s = ±1/2 (2 elements)
p = 3, q = 3, r = 0, s = ±1/2 (2 elements)
p = 3, q = 3, r = +2, s = ±1/2 (2 elements)
4th period:
p = 4; q and r values are the same as with p = 3 (8 total elements)
1
2
3
4
5
6
7
8
9 10 11 12
13 14 15 16 17 18 19 20
b. Elements 2, 4, 12, and 20 all have filled shells and will be least reactive.
c. Draw similarities to the modern periodic table.
XY could be X+Y−, X2+Y2−, or X3+Y3−. Possible ions for each are:
X+ could be elements 1, 3, 5, or 13; Y− could be 11 or 19.
X2+ could be 6 or 14; Y2− could be 10 or 18.
X3+ could be 7 or 15; Y3− could be 9 or 17.
Note: X4+ and Y4− ions probably won’t form.
XY2 will be X2+(Y− )2; See above for possible ions.
X2Y will be (X+)2Y2− See above for possible ions.
XY3 will be X3+(Y− )3; See above for possible ions.
X2Y3 will be (X3+)2(Y2− )3; See above for possible ions.
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
247
d. p = 4, q = 3, r = −2 , s = ±1/2 (2 electrons)
p = 4, q = 3, r = 0, s = ±1/2 (2 electrons)
p = 4, q = 3, r = +2, s = ±1/2 (2 electrons)
A total of 6 electrons can have p = 4 and q = 3.
e.
p = 3, q = 0, r = 0; this is not allowed; q must be odd. Zero electrons can have these
quantum numbers.
f.
p = 6, q = 1, r = 0, s = ±1/2 (2 electrons)
p = 6, q = 3, r = −2, 0, +2; s = ±1/2 (6 electrons)
p = 6, q = 5, r = −4, −2, 0, +2, +4; s = ±1/2 (10 electrons)
Eighteen electrons can have p = 6.
172.
The third ionization energy refers to the following process: E2+(g) → E3+(g) + e− ∆H = I3.
The electron configurations for the 2+ charged ions of Na to Ar are:
Na2+: 1s22s22p5
Mg2+: 1s22s22p6
Al2+:
Si2+:
P2+:
S2+:
Cl2+:
Ar2+:
[Ne]3s1
[Ne]3s2
[Ne]3s23p1
[Ne]3s23p2
[Ne]3s23p3
[Ne]3s23p4
I3 for sodium and magnesium should be extremely large compared with the others because n
= 2 electrons are much more difficult to remove than n = 3 electrons. Between Na2+ and
Mg2+, one would expect to have the same trend as seen with I1(F) versus I1(Ne); these neutral
atoms have identical electron configurations to Na2+ and Mg2+. Therefore, the 1s22s22p5 ion
(Na2+) should have a lower ionization energy than the 1s22s22p6 ion (Mg2+).
The remaining 2+ ions (Al2+ to Ar2+) should follow the same trend as the neutral atoms
having the same electron configurations. The general ionization energy trend predicts an
increase from [Ne]3s1 to [Ne]3s23p4. The exceptions occur between [Ne]3s2 and [Ne]3s23p1
and between [Ne]3s23p3 and [Ne]3s23p4. [Ne]3s23p1 is out of order because of the small
penetrating ability of the 3p electron as compared with the 3s electrons. [Ne]3s23p4 is out of
order because of the extra electron-electron repulsions present when two electrons are paired
in the same orbital. Therefore, the correct ordering for Al2+ to Ar2+ should be Al2+ < P2+ < Si2+
< S2+ < Ar2+ < Cl2+, where P 2+ and Ar2+ are out of line for the same reasons that Al and S are
out of line in the general ionization energy trend for neutral atoms. A qualitative plot of the
third ionization energies for elements Na through Ar follows.
248
CHAPTER 7
ATOMIC STRUCTURE AND PERIODICITY
IE
Na2+
Mg2+
Al2+
Si2+
P2+
S2+
Cl2+
Ar2+
Note: The actual numbers in Table 7.5 support most of this plot. No I3 is given for Na2+, so
you cannot check this. The only deviation from our discussion is I3 for Ar2+ which is greater
than I3 for Cl2+ instead of less than.
173.
The ratios for Mg, Si, P, Cl, and Ar are about the same. However, the ratios for Na, Al, and S
are higher. For Na, the second ionization energy is extremely high because the electron is
taken from n = 2 (the first electron is taken from n = 3). For Al, the first electron requires a
bit less energy than expected by the trend due to the fact it is a 3p electron versus a 3s
electron. For S, the first electron requires a bit less energy than expected by the trend due to
electrons being paired in one of the p orbitals.
174.
Size also decreases going across a period. Sc and Ti along with Y and Zr are adjacent
elements. There are 14 elements (the lanthanides) between La and Hf, making Hf
considerably smaller.
175.
a. Assuming the Bohr model applies to the 1s electron, E1s = −RHZ2/n2 = −RHZ2eff, where
n = 1. Ionization energy = E∞ − E1s = 0 – E1s = RHZ2eff.
2.462 × 10 6 kJ
1 mol
1000 J
= 2.178 × 10−18 J (Zeff)2, Zeff = 43.33
×
×
23
mol
kJ
6.0221 × 10
b. Silver is element 47, so Z = 47 for silver. Our calculated Zeff value is less than 47.
Electrons in other orbitals can penetrate the 1s orbital. Thus a 1s electron can be slightly
shielded from the nucleus by these penetrating electrons, giving a Zeff close to but less
than Z.
176.
None of the noble gases and no subatomic particles had been discovered when Mendeleev
published his periodic table. Thus there was no known element out of place in terms of
reactivity and there was no reason to predict an entire family of elements. Mendeleev ordered
his table by mass; he had no way of knowing there were gaps in atomic numbers (they hadn't
been discovered yet).
CHAPTER 7
177.
m=
ATOMIC STRUCTURE AND PERIODICITY
249
h
6.626 × 10 −34 kg m 2 /s
=
= 6.68 × 10 −26 kg/atom
λv
3.31 × 10 −15 m × (0.0100 × 2.998 × 108 m/s)
6.68 × 10 −26 kg
6.022 × 10 23 atoms
1000 g
×
×
= 40.2 g/mol
atom
mol
1 kg
The element is calcium, Ca.
Integrated Problems
178.
a. ν =
λ=
E
7.52 × 10 −19 J
=
= 1.13 × 1015 s −1
h
6.626 × 10 −34 J s
c
2.998 × 108 m / s
=
= 2.65 × 10 −7 m = 265 nm
ν
1.13 × 1015 s −1
b. Ephoton and λ are inversely related (E = hc/λ). Any wavelength of electromagnetic
radiation less than or equal to 265 nm (λ ≤ 265) will have sufficient energy to eject an
electron. So, yes, 259-nm EMR will eject an electron.
c. This is the electron configuration for copper, Cu, an exception to the expected filling
order.
179.
a. An atom of francium has 87 protons and 87 electrons. Francium is an alkali metal and
forms stable 1+ cations in ionic compounds. This cation would have 86 electrons. The
electron configurations will be:
Fr: [Rn]7s1; Fr+: [Rn] = [Xe]6s24f145d106p6
b. 1.0 oz Fr ×
1 lb
1 kg
1000 g
1 mol Fr
6.02 × 10 23 atoms
×
×
×
×
16 oz 2.205 lb
1 kg
223 g Fr
1 mol Fr
= 7.7 × 1022 atoms Fr
c.
223
Fr is element 87, so it has 223 – 87 = 136 neutrons.
1.67493 × 10 −27 kg 1000 g
136 neutrons ×
×
= 2.27790 × 10 −22 g neutrons
1 neutron
1 kg
180.
a. [Kr]5s24d105p6 = Xe; [Kr]5s24d105p1 = In; [Kr]5s24d105p3 = Sb
From the general radii trend, the increasing size order is Xe < Sb < In.
b. [Ne]3s23p5 = Cl; [Ar]4s23d104p3 = As; [Ar]4s23d104p5 = Br
From the general ionization energy trend, the decreasing ionization energy order is: Cl >
Br > As.
CHAPTER 8
BONDING: GENERAL CONCEPTS
Questions
15.
a. This diagram represents a polar covalent bond as in HCl. In a polar covalent bond,
there is an electron rich region (indicated by the red color) and an electron poor region
(indicated by the blue color). In HCl, the more electronegative Cl atom (on the red side
of the diagram) has a slightly greater ability to attract the bonding electrons than does H
(on the blue side of the diagram), which in turn produces a dipole moment.
b. This diagram represents an ionic bond as in NaCl. Here, the electronegativity differences
between the Na and Cl are so great that the valence electron of sodium is transferred to
the chlorine atom. This results in the formation of a cation, an anion, and an ionic bond.
c. This diagram represents a pure covalent bond as in H2. Both atoms attract the bonding
electrons equally, so there is no bond dipole formed. This is illustrated in the
electrostatic potential diagram as the various red and blue colors are equally distributed
about the molecule. The diagram shows no one region that is red nor one region that is
blue (there is no specific partial negative end and no specific partial positive end), so the
molecule is nonpolar.
16.
In F2 the bonding is pure covalent, with the bonding electrons shared equally between the two
fluorine atoms. In HF, there is also a shared pair of bonding electrons, but the shared pair is
drawn more closely to the fluorine atom. This is called a polar covalent bond as opposed to
the pure covalent bond in F2.
17.
Of the compounds listed, P2O5 is the only compound containing only covalent bonds.
(NH4)2SO4, Ca3(PO4)2, K2O, and KCl are all compounds composed of ions, so they exhibit
ionic bonding. The polyatomic ions in (NH4)2SO4 are NH4+ and SO42−. Covalent bonds exist
between the N and H atoms in NH4+ and between the S and O atoms in SO42−. Therefore,
(NH4)2SO4 contains both ionic and covalent bonds. The same is true for Ca3(PO4)2. The
bonding is ionic between the Ca2+ and PO43− ions and covalent between the P and O atoms in
PO43−. Therefore, (NH4)2SO4 and Ca3(PO4)2 are the compounds with both ionic and covalent
bonds.
18.
Ionic solids are held together by strong electrostatic forces that are omnidirectional.
i.
For electrical conductivity, charged species must be free to move. In ionic solids, the
charged ions are held rigidly in place. Once the forces are disrupted (melting or
dissolution), the ions can move about (conduct).
ii. Melting and boiling disrupts the attractions of the ions for each other. Because these
electrostatic forces are strong, it will take a lot of energy (high temperature) to
accomplish this.
250
CHAPTER 8
BONDING: GENERAL CONCEPTS
251
iii. If we try to bend a piece of material, the ions must slide across each other. For an
ionic solid the following might happen:
strong attraction
strong repulsion
Just as the layers begin to slide, there will be very strong repulsions causing the solid
to snap across a fairly clean plane.
iv. Polar molecules are attracted to ions and can break up the lattice.
These properties and their correlation to chemical forces will be discussed in detail in
Chapters 10 and 11.
19.
Electronegativity increases left to right across the periodic table and decreases from top to
bottom. Hydrogen has an electronegativity value between B and C in the second row and
identical to P in the third row. Going further down the periodic table, H has an electronegativity value between As and Se (row 4) and identical to Te (row 5). It is important to
know where hydrogen fits into the electronegativity trend, especially for rows 2 and 3. If you
know where H fits into the trend, then you can predict bond dipole directions for nonmetals
bonded to hydrogen.
20.
Linear structure (180° bond angle)
S
O
C
Polar; bond dipoles do not cancel.
O
C
O
Nonpolar; bond dipoles cancel.
Trigonal planar structure (120° bond angle)
O
O
C
S
H
H
Polar; bond dipoles do not cancel.
O
O
+ 2 other resonance
structures
Nonpolar; bond dipoles cancel.
Tetrahedral structure (109.5° bond angles)
F
F
C
C
F
H
F
Polar; bond dipoles do not cancel.
F
F
F
Nonpolar; bond dipoles cancel.
252
CHAPTER 8
BONDING: GENERAL CONCEPTS
21.
For ions, concentrate on the number of protons and the number of electrons present. The
species whose nucleus holds the electrons most tightly will be smallest. For example, anions
are larger than the neutral atom. The anion has more electrons held by the same number of
protons in the nucleus. These electrons will not be held as tightly, resulting in a bigger size
for the anion as compared to the neutral atom. For isoelectronic ions, the same number of
electrons are held by different numbers of protons in the various ions. The ion with the most
protons holds the electrons tightest and is smallest in size.
22.
Two other factors that must be considered are the ionization energy needed to produce more
positively charged ions and the electron affinity needed to produce more negatively charged
ions. The favorable lattice energy more than compensates for the unfavorable ionization
energy of the metal and for the unfavorable electron affinity of the nonmetal as long as
electrons are added to or removed from the valence shell. Once the valence shell is empty, the
ionization energy required to remove the next (inner-core) electron is extremely unfavorable;
the same is true for electron affinity when an electron is added to a higher n shell. These two
quantities are so unfavorable after the valence shell is complete that they overshadow the
favorable lattice energy, and the higher charged ionic compounds do not form.
23.
Fossil fuels contain a lot of carbon and hydrogen atoms. Combustion of fossil fuels (reaction
with O2) produces CO2 and H2O. Both these compounds have very strong bonds. Because
stronger product bonds are formed than reactant bonds broken, combustion reactions are very
exothermic.
24.
Statements a and c are true. For statement a, XeF2 has 22 valence electrons, and it is
impossible to satisfy the octet rule for all atoms with this number of electrons. The best Lewis
structure is:
F
Xe
F
For statement c, NO+ has 10 valence electrons, whereas NO− has 12 valence electrons. The
Lewis structures are:
+
N
O
N
O
Because a triple bond is stronger than a double bond, NO+ has a stronger bond.
For statement b, SF4 has five electron pairs around the sulfur in the best Lewis structure; it is
an exception to the octet rule. Because OF4 has the same number of valence electrons as SF4,
OF4 would also have to be an exception to the octet rule. However, row 2 elements such as O
never have more than 8 electrons around them, so OF 4 does not exist. For statement d, two
resonance structures can be drawn for ozone:
O
O
O
O
O
O
When resonance structures can be drawn, the actual bond lengths and strengths are all equal
to each other. Even though each Lewis structure implies the two O−O bonds are different,
this is not the case in real life. In real life, both of the O−O bonds are equivalent. When
resonance structures can be drawn, you can think of the bonding as an average of all of the
resonance structures.
CHAPTER 8
25.
BONDING: GENERAL CONCEPTS
CO2, 4 + 2(6) = 16 valence electrons
--1
0
0
0
O
C
O
O
253
00
+1
+
+1
+
C
O
O
00
--1
C
O
The formal charges are shown above the atoms in the three Lewis structures. The best Lewis
structure for CO2 from a formal charge standpoint is the first structure having each oxygen
double bonded to carbon. This structure has a formal charge of zero on all atoms (which is
preferred). The other two resonance structures have nonzero formal charges on the oxygens,
making them less reasonable. For CO2, we usually ignore the last two resonance structures
and think of the first structure as the true Lewis structure for CO2.
26.
Only statement c is true. The bond dipoles in CF4 and KrF4 are arranged in a manner that they
all cancel each other out, making them nonpolar molecules (CF4 has a tetrahedral molecular
structure, whereas KrF4 has a square planar molecular structure). In SeF4, the bond dipoles in
this see-saw molecule do not cancel each other out, so SeF4 is polar. For statement a, all the
molecules have either a trigonal planar geometry or a trigonal bipyramid geometry, both of
which have 120° bond angles. However, XeCl2 has three lone pairs and two bonded chlorine
atoms around it. XeCl2 has a linear molecular structure with a 180° bond angle. With three
lone pairs, we no longer have a 120° bond angle in XeCl2. For statement b, SO2 has a Vshaped molecular structure with a bond angle of about 120°. CS2 is linear with a 180° bond
angle, and SCl2 is V-shaped but with an approximate 109.5° bond angle. The three compounds do not have the same bond angle. For statement d, central atoms adopt a geometry to
minimize electron repulsions, not maximize them.
Exercises
Chemical Bonds and Electronegativity
27.
Using the periodic table, the general trend for electronegativity is:
(1) Increase as we go from left to right across a period
(2) Decrease as we go down a group
Using these trends, the expected orders are:
a. C < N < O
b. Se < S < Cl
c. Sn < Ge < Si
d. Tl < Ge < S
28.
a. Rb < K < Na
b. Ga < B < O
c. Br < Cl < F
d. S < O < F
29.
The most polar bond will have the greatest difference in electronegativity between the two
atoms. From positions in the periodic table, we would predict:
30.
a. Ge‒F
b. P‒Cl
c. S‒F
d. Ti‒Cl
a. Sn‒H
b. Tl‒Br
c. Si‒O
d. O‒F
254
31.
CHAPTER 8
BONDING: GENERAL CONCEPTS
The general trends in electronegativity used in Exercises 27 and 29 are only rules of
thumb. In this exercise, we use experimental values of electronegativities and can begin to
see several exceptions. The order of EN from Figure 8.3 is:
a. C (2.5) < N (3.0) < O (3.5)
same as predicted
b. Se (2.4) < S (2.5) < Cl (3.0)
same
c. Si = Ge = Sn (1.8)
different
d. Tl (1.8) = Ge (1.8) < S (2.5)
different
Most polar bonds using actual EN values:
a. Si‒F and Ge‒F have equal polarity (Ge‒F predicted).
b. P‒Cl (same as predicted)
c. S‒F (same as predicted)
32.
d. Ti‒Cl (same as predicted)
The order of EN from Figure 8.3 is:
a. Rb (0.8) = K (0.8) < Na (0.9), different
b. Ga (1.6) < B (2.0) < O (3.5), same
c. Br (2.8) < Cl (3.0) < F (4.0), same
d. S (2.5) < O (3.5) < F (4.0), same
Most polar bonds using actual EN values:
a. C‒H most polar (Sn‒H predicted)
b. Al‒Br most polar (Tl‒Br predicted).
c. Si‒O (same as predicted).
d. Each bond has the same polarity, but the bond dipoles point in opposite directions.
Oxygen is the positive end in the O‒F bond dipole, and oxygen is the negative end in the
O‒Cl bond dipole (O‒F predicted).
33.
Use the electronegativity trend to predict the partial negative end and the partial positive end
of the bond dipole (if there is one). To do this, you need to remember that H has electronegativity between B and C and identical to P. Answers b, d, and e are incorrect. For d (Br2),
the bond between two Br atoms will be a pure covalent bond, where there is equal sharing of
the bonding electrons, and no dipole moment exists. For b and e, the bond polarities are
reversed. In Cl−I, the more electronegative Cl atom will be the partial negative end of the
bond dipole, with I having the partial positive end. In O−P, the more electronegative oxygen
will be the partial negative end of the bond dipole, with P having the partial positive end. In
the following, we used arrows to indicate the bond dipole. The arrow always points to the
partial negative end of a bond dipole (which always is the most electronegative atom in the
bond).
Cl
I
O
P
CHAPTER 8
34.
BONDING: GENERAL CONCEPTS
See Exercise 33 for a discussion on bond dipoles. We will use arrows to indicate the bond
dipoles. The arrow always points to the partial negative end of the bond dipole, which will
always be to the more electronegative atom. The tail of the arrow indicates the partial positive
end of the bond dipole.
a.
35.
36.
255
C
O
b. P−H is a pure covalent (nonpolar) bond because
P and H have identical electronegativities.
c.
H
Cl
d.
e.
Se
S
Br
Te
The actual electronegativity difference between Se and S is so small that
this bond is probably best characterized as a pure covalent bond having
no bond dipole.
Bonding between a metal and a nonmetal is generally ionic. Bonding between two nonmetals
is covalent, and in general, the bonding between two different nonmetals is usually polar
covalent. When two different nonmetals have very similar electronegativities, the bonding is
pure covalent or just covalent.
a. ionic
b. covalent
c. polar covalent
d. ionic
e. polar covalent
f.
covalent
The possible ionic bonds that can form are between the metal Cs and the nonmetals P, O, and
H. These ionic compounds are Cs3P, Cs2O, and CsH. The bonding between the various
nonmetals will be covalent. P4, O2, and H2 are all pure covalent (or just covalent) with equal
sharing of the bonding electrons. P−H will also be a covalent bond because P and H have
identical electronegativities. The other possible covalent bonds that can form will all be polar
covalent because the nonmetals involved in the bonds all have intermediate differences in
electronegativities. The possible polar covalent bonds are P−O and O−H.
Note: The bonding among cesium atoms is called metallic. This type of bonding between
metals will be discussed in Chapter 10.
37.
Electronegativity values increase from left to right across the periodic table. The order of
electronegativities for the atoms from smallest to largest electronegativity will be H = P < C <
N < O < F. The most polar bond will be F‒H since it will have the largest difference in
electronegativities, and the least polar bond will be P‒H since it will have the smallest
difference in electronegativities (ΔEN = 0). The order of the bonds in decreasing polarity will
be F‒H > O‒H > N‒H > C‒H > P‒H.
38.
Ionic character is proportional to the difference in electronegativity values between the two
elements forming the bond. Using the trend in electronegativity, the order will be:
Br‒Br < N‒O < C‒F < Ca‒O < K‒F
least
most
ionic character
ionic character
256
CHAPTER 8
BONDING: GENERAL CONCEPTS
Note that Br‒Br, N‒O, and C‒F bonds are all covalent bonds since the elements are all nonmetals. The Ca‒O and K‒F bonds are ionic, as is generally the case when a metal forms a
bond with a nonmetal.
39.
40.
A permanent dipole moment exists in a molecule if the molecule has one specific area with a
partial negative end (a red end in an electrostatic potential diagram) and a different specific
region with a partial positive end (a blue end in an electrostatic potential diagram). If the
blue and red colors are equally distributed in the electrostatic potential diagrams, then no
permanent dipole exists.
a. Has a permanent dipole.
b. Has no permanent dipole.
c. Has no permanent dipole.
d. Has a permanent dipole.
e. Has no permanent dipole.
f.
Has no permanent dipole.
a. H2O; both H2O and NH3 have permanent dipole moments in part due to the polar O−H
and N−H bonds. But because oxygen is more electronegative than nitrogen, one would
expect H2O to have a slightly greater dipole moment. This diagram has the more intense
red color on one end and the more intense blue color at the other end indicating a larger
dipole moment.
b. NH3; this diagram is for a polar molecule, but the colors are not as intense as the diagram
in part a. Hence, this diagram is for a molecule which is not as polar as H2O. Since N is
less electronegative than O, NH3 will not be as polar as H2O.
c. CH4; this diagram has no one specific red region and has four blue regions arranged
symmetrically about the molecule. This diagram is for a molecule which has no dipole
moment. This is only true for CH4. The C‒H bonds are at best, slightly polar because
carbon and hydrogen have similar electronegativity values. In addition, the slightly polar
C‒H bond dipoles are arranged about carbon so that they cancel each other out, making
CH4 a nonpolar molecule. See Example 8.2.
Ions and Ionic Compounds
41.
Al3+: [He]2s22p6; Ba2+: [Kr]5s24d105p6; Se2−: [Ar]4s23d104p6 ; I−: [Kr]5s24d105p6
42.
Te2−: [Kr]5s24d105p6; Cl−: [Ne]3s23p6; Sr2+: [Ar]4s23d104p6; Li+: 1s2
43.
a. Li+ and N3− are the expected ions. The formula of the compound would be Li3N (lithium
nitride).
b. Ga3+ and O2−; Ga2O3, gallium(III) oxide or gallium oxide
44.
c. Rb+ and Cl−; RbCl, rubidium chloride
d. Ba2+ and S2−; BaS, barium sulfide
a. Al3+ and Cl−; AlCl3, aluminum chloride
b. Na+ and O2−; Na2O, sodium oxide
CHAPTER 8
BONDING: GENERAL CONCEPTS
c. Sr2+ and F−; SrF2, strontium fluoride
45.
257
d. Ca2+ and S2−; CaS, calcium sulfide
a. Mg2+: 1s22s22p6; K+: 1s22s22p63s23p6; Al3+: 1s22s22p6
b. N3−, O2−, and F−: 1s22s22p6; Te2-: [Kr]5s24d105p6
46.
a. Sr2+: [Ar]4s23d104p6; Cs+: [Kr]5s24d105p6; In+: [Kr]5s24d10; Pb2+: [Xe]6s24f145d10
b. P3− and S2−: [Ne]3s23p6; Br−: [Ar]4s23d104p6
47.
a. Sc3+: [Ar]
b. Te2−: [Xe]
c. Ce4+: [Xe] and Ti4+: [Ar]
d. Ba2+: [Xe]
All these ions have the noble gas electron configuration shown in brackets.
48.
a. Cs2S is composed of Cs+ and S2−. Cs+ has the same electron configuration as Xe, and S2−
has the same configuration as Ar.
b. SrF2; Sr2+ has the Kr electron configuration, and F− has the Ne configuration.
c. Ca3N2; Ca2+ has the Ar electron configuration, and N3− has the Ne configuration.
d. AlBr3; Al3+ has the Ne electron configuration, and Br− has the Kr configuration.
49.
a. Na+ has 10 electrons. F−, O2−, and N3− are some possible anions also having 10 electrons.
b. Ca2+ has 18 electrons. Cl−, S2−, and P3− also have 18 electrons.
c. Al3+ has 10 electrons. F−, O2−, and N3− also have 10 electrons.
d. Rb+ has 36 electrons. Br−, Se2−, and As3− also have 36 electrons.
50.
a. Ne has 10 electrons. AlN, MgF2, and Na2O are some possible ionic compounds where
each ion has 10 electrons.
b. CaS, K3P, and KCl are some examples where each ion is isoelectronic with Ar; i.e., each
ion has 18 electrons.
c. Each ion in Sr3As2, SrBr2, and Rb2Se is isoelectronic with Kr.
d. Each ion in BaTe and CsI is isoelectronic with Xe.
51.
Neon has 10 electrons; there are many possible ions with 10 electrons. Some are N3−, O2−,
F−, Na+, Mg2+, and Al3+. In terms of size, the ion with the most protons will hold the
electrons the tightest and will be the smallest. The largest ion will be the ion with the fewest
protons. The size trend is:
Al3+ < Mg2+ < Na+ < F− < O2− < N3−
smallest
largest
258
52.
CHAPTER 8
All these ions have 18 e−; the smallest ion (Sc3+) has the most protons attracting the 18 e−,
and the largest ion has the fewest protons (S2−). The order in terms of increasing size is Sc3+
< Ca2+ < K+ < Cl− < S2−. In terms of the atom size indicated in the question:
K+
53.
BONDING: GENERAL CONCEPTS
Ca2+
Sc3+
S2-
Cl
c. O2− > O− > O
a. Cu > Cu+ > Cu2+
b. Pt2+ > Pd2+ > Ni2+
d. La3+ > Eu3+ > Gd3+ > Yb3+
e. Te2− > I− > Cs+ > Ba2+ > La3+
For answer a, as electrons are removed from an atom, size decreases. Answers b and d follow
the radius trend. For answer c, as electrons are added to an atom, size increases. Answer e
follows the trend for an isoelectronic series; i.e., the smallest ion has the most protons.
54.
55.
56.
57.
a. V > V2+ > V3+ > V5+
b. Cs+ > Rb+ > K+ > Na+
d. P3− > P2− > P− > P
e. Te2− > Se2− > S2− > O2−
c. Te2− > I− > Cs+ > Ba2+
Lattice energy is proportional to −Q1Q2/r, where Q is the charge of the ions and r is the
distance between the centers of the ions. The more negative the lattice energy, the more
stable the ionic compound. So greater charged ions as well as smaller sized ions lead to more
negative lattice energy values and more stable ionic compounds.
a. NaCl; Na+ is smaller than K+.
b. LiF; F− is smaller than Cl−.
c. MgO; O2− has a greater charge than OH-.
d. Fe(OH)3; Fe3+ has a greater charge than
Fe2+.
e. Na2O; O2− has a greater charge than Cl−.
f. MgO; both ions are smaller in MgO.
a. LiF; Li+ is smaller than Cs+.
b. NaBr; Br- is smaller than I−.
c. BaO; O2− has a greater charge than Cl-.
d. CaSO4; Ca2+ has a greater charge than Na+.
e. K2O; O2− has a greater charge than F-.
f. Li2O; both ions are smaller in Li2O.
K(s) → K(g)
K(g) → K+(g) + e−
1/2 Cl2(g) → Cl(g)
Cl(g) + e−→ Cl−(g)
K+(g) + Cl−(g) → KCl(s)
K(s) + 1/2 Cl2(g) → KCl(s)
ΔH = 90. kJ (sublimation)
ΔH = 419 kJ (ionization energy)
ΔH = 239/2 kJ (bond energy)
ΔH = −349 kJ (electron affinity)
ΔH = −690. kJ (lattice energy)
ΔH of = −411 kJ/mol
CHAPTER 8
58.
BONDING: GENERAL CONCEPTS
Mg(s) → Mg(g)
Mg(g) → Mg+(g) + e−
Mg+(g) → Mg2+(g) + e−
F2(g) → 2 F(g)
2 F(g) + 2 e− → 2 F−(g)
Mg2+(g) + 2 F−(g) → MgF2(s)
Mg(s) + F2(g) → MgF2(s)
59.
60.
ΔH = 150. kJ
ΔH = 735 kJ
ΔH = 1445 kJ
ΔH = 154 kJ
ΔH = 2(−328) kJ
ΔH = −2913 kJ
259
(sublimation)
(IE1)
(IE2)
(BE)
(EA)
(LE)
ΔH of = −1085 kJ/mol
From the data given, it takes less energy to produce Mg+(g) + O−(g) than to produce
Mg2+(g) + O2−(g). However, the lattice energy for Mg2+O2− will be much more exothermic
than that for Mg+O− due to the greater charges in Mg2+O2−. The favorable lattice energy term
dominates, and Mg2+O2− forms.
Na(g) → Na+(g) + e−
F(g) + e− → F−(g)
Na(g) + F(g) → Na+(g) + F−(g)
ΔH = IE = 495 kJ (Table 7.5)
ΔH = EA = −327.8 kJ (Table 7.7)
ΔH = 167 kJ
The described process is endothermic. What we haven’t accounted for is the extremely favorable lattice energy. Here, the lattice energy is a large negative (exothermic) value, making the
overall formation of NaF a favorable exothermic process.
61.
Use Figure 8.11 as a template for this problem.
Li(s) →
Li(g) →
1/2 I2(g) →
I(g) + e− →
+
Li (g) + I−(g) →
Li(g)
Li+(g) + e−
I(g)
I−(g)
LiI(s)
Li(s) + 1/2 I2(g) → LiI(s)
ΔHsub = ?
ΔH = 520. kJ
ΔH = 151/2 kJ
ΔH = −295 kJ
ΔH = −753 kJ
ΔH = −292 kJ
ΔHsub + 520. + 151/2 − 295 − 753 = −292, ΔHsub = 161 kJ
62.
Let us look at the complete cycle for Na2S.
2 Na(s) → 2 Na(g)
2 Na(g) → 2 Na+(g) + 2 e−
S(s) → S(g)
S(g) + e− → S−(g)
S−(g) + e− → S2−(g)
2 Na+(g) + S2−(g) → Na2S(s)
2 Na(s) + S(s) → Na2S(s)
2ΔHsub, Na = 2(109) kJ
2IE = 2(495) kJ
ΔHsub, S = 277 kJ
EA1 = −200. kJ
EA2 = ?
LE = −2203 kJ
ΔH of = −365 kJ
ΔH of = 2ΔH sub , Na + 2IE + ΔHsub, S + EA1 + EA2 + LE, −365 = −918 + EA2, EA2 = 553 kJ
260
CHAPTER 8
BONDING: GENERAL CONCEPTS
For each salt: ΔH of = 2ΔHsub, M + 2IE + 277 − 200. + LE + EA2
K2S: −381 = 2(90.) + 2(419) + 277 − 200. − 2052 + EA2, EA2 = 576 kJ
Rb2S: −361 = 2(82) + 2(409) + 277 − 200. − 1949 + EA2, EA2 = 529 kJ
Cs2S: −360. = 2(78) + 2(382) + 277 − 200. − 1850. + EA2, EA2 = 493 kJ
We get values from 493 to 576 kJ.
The mean value is:
540 ±50 kJ.
63.
553 + 576 + 529 + 493
= 538 kJ. We can represent the results as EA2 =
4
Ca2+ has a greater charge than Na+, and Se2− is smaller than Te2−. The effect of charge on the
lattice energy is greater than the effect of size. We expect the trend from most exothermic
lattice energy to least exothermic to be:
CaSe > CaTe > Na2Se > Na2Te
(−2862) (−2721) (−2130) (−2095)
64.
This is what we observe.
Lattice energy is proportional to the charge of the cation times the charge of the anion Q1Q2.
Compound
Q1Q2
Lattice Energy
FeCl2
(+2)( −1) = −2
−2631 kJ/mol
FeCl3
(+3)( −1) = −3
−5359 kJ/mol
Fe2O3
(+3)( −2) = −6
−14,744 kJ/mol
Bond Energies
65.
a.
H + Cl
H
Cl
Bonds broken:
2H
Cl
Bonds formed:
1 H‒H (432 kJ/mol)
1 Cl‒Cl (239 kJ/mol)
2 H‒Cl (427 kJ/mol)
ΔH = ΣDbroken − ΣDformed, ΔH = 432 kJ + 239 kJ − 2(427) kJ = −183 kJ
b.
N
N+ 3 H
H
2H
N
H
H
Bonds broken:
1 N≡N (941 kJ/mol)
3 H‒H (432 kJ/mol)
Bonds formed:
6 N‒H (391 kJ/mol)
ΔH = 941 kJ + 3(432) kJ − 6(391) kJ = −109 kJ
CHAPTER 8
66.
BONDING: GENERAL CONCEPTS
261
Sometimes some of the bonds remain the same between reactants and products. To save
time, only break and form bonds that are involved in the reaction.
H
H
C
N
H
H
a.
H
C
N+2 H
H
H
Bonds broken:
Bonds formed:
1 C≡N (891 kJ/mol)
2 H−H (432 kJ/mol)
1 C−N (305 kJ/mol)
2 C−H (413 kJ/mol)
2 N−H (391 kJ/mol)
ΔH = 891 kJ + 2(432 kJ) − [305 kJ + 2(413 kJ) + 2(391 kJ)] = −158 kJ
H
b.
H
N
+2 F
N
H
F + N
4 H
F
N
H
Bonds broken:
Bonds formed:
1 N−N (160. kJ/mol)
4 N−H (391 kJ/mol)
2 F−F (154 kJ/mol)
4 H−F (565 kJ/mol)
1 N ≡ N (941 kJ/mol)
ΔH = 160. kJ + 4(391 kJ) + 2(154 kJ) − [4(565 kJ) + 941 kJ] = −1169 kJ
H
H
67.
H
C N
C
H
H
Bonds broken: 1 C‒N (305 kJ/mol)
C C N
H
Bonds formed: 1 C‒C (347 kJ/mol)
ΔH = ΣDbroken − ΣDformed, ΔH = 305 − 347 = −42 kJ
Note: Sometimes some of the bonds remain the same between reactants and products.
To save time, only break and form bonds that are involved in the reaction.
68.
H
H
C O H + C O
H
H
H
O
C
C O H
H
262
CHAPTER 8
Bonds broken:
BONDING: GENERAL CONCEPTS
Bonds formed:
1 C≡O (1072 kJ/mol)
1 C‒O (358 kJ/mol)
1 C‒C (347 kJ/mol)
1 C=O (745 kJ/mol)
1 C‒O (358 kJ/mol)
ΔH = 1072 + 358 − [347 + 745 + 358] = −20. kJ
F
69.
H
+ 3F
H +
S
F
F
S
+ 2H
F +
F
Bonds broken:
Bonds formed:
2 S−H (347 kJ/mol)
3 F−F (154 kJ/mol)
4 S−F (327 kJ/mol)
2 H−F (565 kJ/mol)
∆H = 2(347) + 3(154) − [4(327) + 2(565)] = −1282 kJ
70.
H
H
C
H + H
O
C
H
O + 3H
H
H
Bonds broken:
4 C−H (413 kJ/mol)
2 O−H (467 kJ/mol)
Bonds formed:
1 C≡O (1072 kJ/mol)
3 H−H (432 kJ/mol)
∆H = 4(413) + 2(467) – [1072 + 3(432)] = 218 kJ
71.
H−C≡C−H + 5/2 O=O → 2 O=C=O + H−O−H
Bonds broken:
2 C−H (413 kJ/mol)
1 C≡C (839 kJ/mol)
5/2 O=O (495 kJ/mol)
Bonds formed:
2 × 2 C=O (799 kJ/mol)
2 O−H (467 kJ/mol)
∆H = 2(413 kJ) + 839 kJ + 5/2 (495 kJ) – [4(799 kJ) + 2(467 kJ)] = −1228 kJ
F
CHAPTER 8
72.
CH4
BONDING: GENERAL CONCEPTS
+
263
2 O=O → O=C=O + 2 H−O−H
Bonds broken:
Bonds formed:
4 C−H (413 kJ/mol)
2 O=O (495 kJ/mol)
2 C=O (799 kJ/mol)
2 × 2 O−H (467 kJ/mol)
∆H = 4(413 kJ) + 2(495 kJ) – [2(799 kJ) + 4(467 kJ)] = −824 kJ
73.
H
H
C
H
+ F
C
F
H
H
F
F
C
C
H
H
Bonds broken:
H
∆H = -549 kJ
Bonds formed:
1 C=C (614 kJ/mol)
1 F−F (154 kJ/mol)
1 C‒C (347 kJ/mol)
2 C‒F (DCF = C‒F bond energy)
ΔH = −549 kJ = 614 kJ + 154 kJ − [347 kJ + 2DCF], 2DCF = 970., DCF = 485 kJ/mol
74.
Let x = bond energy for A2, so 2x = bond energy for AB.
∆H = −285 kJ = x + 432 kJ – [2(2x)], 3x = 717, x = 239 kJ/mol
The bond energy for A2 is 239 kJ/mol.
75.
a. ΔH° = 2 ∆H fo, HCl = 2 mol(−92 kJ/mol) = −184 kJ (−183 kJ from bond energies)
b. ΔH° = 2 ∆H fo, NH = 2 mol(–46 kJ/mol) = −92 kJ (−109 kJ from bond energies)
3
Comparing the values for each reaction, bond energies seem to give a reasonably good
estimate for the enthalpy change of a reaction. The estimate is especially good for gas phase
reactions.
76.
CH3OH(g) + CO(g) → CH3COOH(l)
ΔH° = −484 kJ − [(−201 kJ) + (−110.5 kJ)] = −173 kJ
Using bond energies, ΔH = −20. kJ. For this reaction, bond energies give a much poorer
estimate for ΔH as compared with gas-phase reactions (see Exercise 75). The major reason
for the large discrepancy is that not all species are gases in Exercise 68. Bond energies do not
account for the energy changes that occur when liquids and solids form instead of gases.
These energy changes are due to intermolecular forces and will be discussed in Chapter 10.
77.
a. Using SF4 data: SF4(g) → S(g) + 4 F(g)
ΔH° = 4DSF = 278.8 + 4 (79.0) − (−775) = 1370. kJ
DSF =
1370. kJ
= 342.5 kJ/mol = S−F bond energy
4 mol SF bonds
264
CHAPTER 8
BONDING: GENERAL CONCEPTS
Using SF6 data: SF6(g) → S(g) + 6 F(g)
ΔH° = 6DSF = 278.8 + 6 (79.0) − (−1209) = 1962 kJ
DSF =
1962 kJ
= 327.0 kJ/mol = S−F bond energy
6 mol
b. The S‒F bond energy in Table 8.4 is 327 kJ/mol. The value in the table was based on the
S−F bond in SF6.
c. S(g) and F(g) are not the most stable forms of the elements at 25°C and 1 atm. The most
stable forms are S8(s) and F2(g); ∆H °f = 0 for these two species.
78.
NH3(g) → N(g) + 3 H(g) ΔH° = 3DNH = 472.7 + 3(216.0) − (−46.1) = 1166.8 kJ
DNH =
1166.8 kJ
= 388.93 kJ/mol ≈ 389 kJ/mol
3 mol NH bonds
Dcalc = 389 kJ/mol as compared with 391 kJ/mol in Table 8.4. There is good agreement.
79.
H
H
N
H
2 N(g) + 4 H(g)
N
ΔH = DN‒N + 4DN‒H = DN‒N + 4(388.9)
H
ΔH° = 2 ∆H fo, N + 4 ∆H fo, H − ∆H fo, N
2H 4
= 2(472.7 kJ) + 4(216.0 kJ) − 95.4 kJ
ΔH° = 1714.0 kJ = DN‒N + 4(388.9)
DN‒N = 158.4 kJ/mol (versus 160. kJ/mol in Table 8.4)
80.
1/2 N2(g) + 1/2 O2(g) → NO(g)
Bonds broken:
1/2 N≡N (941 kJ/mol)
1/2 O=O (495 kJ/mol)
∆H = 90. kJ
Bonds formed:
1 NO (DNO = NO bond energy)
∆H = 90. kJ = 1/2(941) + 1/2(495) – (DNO), DNO = 628 kJ/mol
From this data, the calculated NO bond energy is 628 kJ/mol.
CHAPTER 8
BONDING: GENERAL CONCEPTS
265
Lewis Structures and Resonance
81.
Drawing Lewis structures is mostly trial and error. However, the first two steps are always
the same. These steps are (1) count the valence electrons available in the molecule/ion, and
(2) attach all atoms to each other with single bonds (called the skeletal structure). Unless
noted otherwise, the atom listed first is assumed to be the atom in the middle, called the
central atom, and all other atoms in the formula are attached to this atom. The most notable
exceptions to the rule are formulas that begin with H, e.g., H2O, H2CO, etc. Hydrogen can
never be a central atom since this would require H to have more than two electrons. In these
compounds, the atom listed second is assumed to be the central atom.
After counting valence electrons and drawing the skeletal structure, the rest is trial and error.
We place the remaining electrons around the various atoms in an attempt to satisfy the octet
rule (or duet rule for H). Keep in mind that practice makes perfect. After practicing, you can
(and will) become very adept at drawing Lewis structures.
a. F2 has 2(7) = 14 valence electrons.
FF
FF
FF
Skeeleettaal
ssttrruuccttuurree
b. O2 has 2(6) = 12 valence electrons.
FF
O
S kel etal
ssttrruuccttuurree
wiss
Leew
L
ssttrruuccttuurree
c. CO has 4 + 6 = 10 valence electrons.
C
C
O
Skeletal
structure
O
O
O
Lewis
ssttrruuccttuurree
d. CH4 has 4 + 4(1) = 8 valence electrons.
H
H
O
Lewis
structure
H
C
H
H
Lewis
structure
Skeletal
structure
H
N
H
H
N
H
H
Skeletal
structure
Lewis
structure
H
f.
H
H
H
H
e. NH3 has 5 + 3(1) = 8 valence electrons.
C
H2O has 2(1) + 6 = 8 valence electrons.
O
Skeletal
structure
H
H
O
Lewis
structure
H
266
CHAPTER 8
BONDING: GENERAL CONCEPTS
g. HF has 1 + 7 = 8 valence electrons.
H
H
F
Skeletal
structure
82.
F
Lewis
structure
a. H2CO has 2(1) + 4 + 6 = 12 valence
electrons.
O
H
C
b. CO2 has 4 + 2(6) = 16 valence electrons.
O
H
H
Skeletal
structure
C
O
H
C
Skeletal
structure
O
O
C
O
Lewis
structure
Lewis
structure
c. HCN has 1 + 4 + 5 = 10 valence electrons.
H
C
N
H
Skeletal
structure
83.
C
N
Lewis
structure
Drawing Lewis structures is mostly trial and error. However, the first two steps are always
the same. These steps are (1) count the valence electrons available in the molecule/ion, and
(2) attach all atoms to each other with single bonds (called the skeletal structure). Unless
noted otherwise, the atom listed first is assumed to be the atom in the middle, called the
central atom, and all other atoms in the formula are attached to this atom. The most notable
exceptions to the rule are formulas that begin with H, e.g., H2O, H2CO, etc. Hydrogen can
never be a central atom since this would require H to have more than two electrons. In these
compounds, the atom listed second is assumed to be the central atom.
After counting valence electrons and drawing the skeletal structure, the rest is trial and error.
We place the remaining electrons around the various atoms in an attempt to satisfy the octet
rule (or duet rule for H).
a. CCl4 has 4 + 4(7) = 32 valence
electrons.
Cl
Cl
Cl
C
Cl
Cl
C
Cl
Cl
Skeletal
structure
Lewis
structure
b. NCl3 has 5 + 3(7) = 26 valence
electrons.
N
Cl
Cl
N
Cl
Cl
Cl
Skeletal
structure
Cl
Lewis
structure
Cl
CHAPTER 8
BONDING: GENERAL CONCEPTS
c. SeCl2 has 6 + 2(7) = 20 valence
electrons.
Se
Cl
Skeletal
structure
84.
d. ICl has 7 + 7 = 14 valence electrons.
Se
Cl
Cl
267
I
Cl
I
Cl
Skeletal
structure
Lewis
structure
Cl
Lewis
structure
a. POCl3 has 5 + 6 + 3(7) = 32 valence electrons.
O
Cl
O
P
Cl
Cl
P
Cl
Cl
Cl
Skeletal
structure
Lewis
structure
Note: This structure uses all 32 e− while
satisfying the octet rule for all atoms. This is
a valid Lewis structure.
SO42− has 6 + 4(6) + 2 = 32 valence electrons.
2-
O
O
S
Note: A negatively charged ion will have
additional electrons to those that come from
the valence shell of the atoms. The magnitude of the negative charge indicates the
number of extra electrons to add in.
O
O
XeO4, 8 + 4(6) = 32 e−
PO43−, 5 + 4(6) + 3 = 32 e−
3-
O
O
O
Xe
O
O
O
P
O
O
ClO4− has 7 + 4(6) + 1 = 32 valence electrons
-
O
O
Cl
O
O
Note: All of these species have the same number of atoms and the same number of
valence electrons. They also have the same Lewis structure.
268
CHAPTER 8
SO32−, 6 + 3(6) + 2 = 26 e−
b. NF3 has 5 + 3(7) = 26 valence electrons.
F
N
F
F
F
N
BONDING: GENERAL CONCEPTS
2-
F
O
F
Skeletal
structure
S
O
O
Lewis
structure
PO33−, 5 + 3(6) + 3 = 26 e−
3P
O
O
ClO3−, 7 + 3(6) + 1 = 26 e−
O
O
Cl
O
O
Note: Species with the same number of atoms and valence electrons have similar Lewis
structures.
c. ClO2− has 7 + 2(6) + 1 = 20 valence
O Cl
O Cl
O
Skeletal structure
Lewis structure
SCl2, 6 + 2(7) = 20 e−
Cl
S
O
PCl2−, 5 + 2(7) + 1 = 20 e−
P
Cl
Cl
Cl
-
Note: Species with the same number of atoms and valence electrons have similar Lewis
structures.
d. Molecules ions that have the same number of valence electrons and the same number of
atoms will have similar Lewis structures.
85.
BeH2, 2 + 2(1) = 4 valence electrons
BH3, 3 + 3(1) = 6 valence electrons
H
H
Be
H
B
H
H
CHAPTER 8
86.
BONDING: GENERAL CONCEPTS
a. NO2, 5 + 2(6) = 17 e−
N2O4, 2(5) + 4(6) = 34 e−
O
O
N
N
O
O
N
O
269
O
Plus others
Plus other resonance structures
b. BH3, 3 + 3(1) = 6 e−
NH3, 5 + 3(1) = 8 e−
H
N
H
H
B
H
H
H
BH3NH3, 6 + 8 = 14 e−
H
H
H
B
H
N
H
H
In reaction a, NO2 has an odd number of electrons, so it is impossible to satisfy the octet
rule. By dimerizing to form N2O4, the odd electron on two NO2 molecules can pair up,
giving a species whose Lewis structure can satisfy the octet rule. In general, odd-electron
species are very reactive. In reaction b, BH3 is electron-deficient. Boron has only six
electrons around it. By forming BH3NH3, the boron atom satisfies the octet rule by
accepting a lone pair of electrons from NH3 to form a fourth bond.
87.
PF5, 5 +5(7) = 40 valence electrons
SF4, 6 + 4(7) = 34 e−
F
F
F
F
F
F
F
P
S
F
F
ClF3, 7 + 3(7) = 28 e−
Br3−, 3(7) + 1 = 22 e−
F
Br
Cl
F
F
Br
Br
270
CHAPTER 8
BONDING: GENERAL CONCEPTS
Row 3 and heavier nonmetals can have more than 8 electrons around them when they have to.
Row 3 and heavier elements have empty d orbitals that are close in energy to valence s and p
orbitals. These empty d orbitals can accept extra electrons.
For example, P in PF5 has its five valence electrons in the 3s and 3p orbitals. These s and p
orbitals have room for three more electrons, and if it has to, P can use the empty 3d orbitals
for any electrons above 8.
88.
SF6, 6 + 6(7) = 48 e−
F
F
F
S
F
F
F
ClF5, 7 + 5(7) = 42 e−
F
F
F
Cl
F
F
XeF 4, 8 + 4(7) = 36 e−
F
F
Xe
F
89.
F
a. NO2− has 5 + 2(6) + 1 = 18 valence electrons. The skeletal structure is O‒N‒O.
To get an octet about the nitrogen and only use 18 e- , we must form a double bond to one
of the oxygen atoms.
O N
O
O N
O
Because there is no reason to have the double bond to a particular oxygen atom, we can
draw two resonance structures. Each Lewis structure uses the correct number of electrons
and satisfies the octet rule, so each is a valid Lewis structure. Resonance structures occur
when you have multiple bonds that can be in various positions. We say the actual
structure is an average of these two resonance structures.
NO3− has 5 + 3(6) + 1 = 24 valence electrons. We can draw three resonance structures for
NO3−, with the double bond rotating among the three oxygen atoms.
O
O
N
O
O
N
O
O
N
O
O
O
CHAPTER 8
BONDING: GENERAL CONCEPTS
271
N2O4 has 2(5) + 4(6) = 34 valence electrons. We can draw four resonance structures for
N2O4.
O
N
O
O
O
N
N
O
O
O
O
N
N
O
O
O
O
O
N
N
O
O
N
O
b. OCN− has 6 + 4 + 5 + 1 = 16 valence electrons. We can draw three resonance structures
for OCN−.
O C N
O C N
O C N
SCN− has 6 + 4 + 5 + 1 = 16 valence electrons. Three resonance structures can be drawn.
S C N
S
C N
S C N
N3− has 3(5) + 1 = 16 valence electrons. As with OCN− and SCN−, three different
resonance structures can be drawn.
N N
N N N
90.
N
N N
Ozone: O3 has 3(6) = 18 valence electrons.
O
O
O
O
O
O
Sulfur dioxide: SO2 has 6 + 2(6) = 18 valence electrons.
O
S
O
O
S
O
N
272
CHAPTER 8
BONDING: GENERAL CONCEPTS
Sulfur trioxide: SO3 has 6 + 3(6) = 24 valence electrons.
O
O
O
S
S
S
O
91.
O
O
O
O
Benzene has 6(4) + 6(1) = 30 valence electrons. Two resonance structures can be drawn for
benzene. The actual structure of benzene is an average of these two resonance structures; i.e.,
all carbon-carbon bonds are equivalent with a bond length and bond strength somewhere
between a single and a double bond.
H
H
H
H
C
C
C
H
C
C
C
H
C
C
C
H
H
H
H
Borazine (B3N3H6) has 3(3) + 3(5) + 6(1) = 30 valence electrons. The possible resonance
structures are similar to those of benzene in Exercise 91.
H
H
H
B
H
N
H
H
B
N
N
B
B
H
H
H
B
N
N
B
N
H
H
H
93.
H
C
C
C
H
92.
O
We will use a hexagon to represent the six-member carbon ring, and we will omit the four
hydrogen atoms and the three lone pairs of electrons on each chlorine. If no resonance
existed, we could draw four different molecules:
Cl
Cl
Cl
Cl
Cl
Cl
Cl
Cl
CHAPTER 8
BONDING: GENERAL CONCEPTS
273
If the double bonds in the benzene ring exhibit resonance, then we can draw only three
different dichlorobenzenes. The circle in the hexagon represents the delocalization of the
three double bonds in the benzene ring (see Exercise 91).
Cl
Cl
Cl
Cl
Cl
Cl
With resonance, all carbon-carbon bonds are equivalent. We can’t distinguish between a
single and double bond between adjacent carbons that have a chlorine attached. That only
three isomers are observed supports the concept of resonance.
CO32− has 4 + 3(6) + 2 = 24 valence electrons.
94.
2-
O
C
O
2-
O
C
C
O
2-
O
O
O
O
O
Three resonance structures can be drawn for CO32−. The actual structure for CO32− is an
average of these three resonance structures. That is, the three C‒O bond lengths are all
equivalent, with a length somewhere between a single and a double bond. The actual bond
length of 136 pm is consistent with this resonance view of CO32−.
95.
CH3NCO has 4 + 3(1) + 5 + 4 + 6 = 22 valence electrons. Three resonance structures can be
drawn for methyl isocyanate.
H
H
C
H
N
C
O
H
H
96.
C
H
N
C
O
H
H
C
N
C
H
PAN (H3C2NO5) has 3(1) + 2(4) + 5 + 5(6) = 46 valence electrons.
H
H
O
C
C
H
O
O
O
N
O
This is the skeletal structure with complete octets
about oxygen atoms (46 electrons used).
O
274
CHAPTER 8
BONDING: GENERAL CONCEPTS
This structure has used all 46 electrons, but there are only six electrons around one of the
carbon atoms and the nitrogen atom. Two unshared pairs must become shared; i.e., we must
form double bonds. The three possible resonance structures for PAN are:
H
H
O
C
C
O
O
O
N
H
O
H
H
H
O
C
C
O
C
C
O
O
O
N
O
H
O
O
O
(last form not important)
N
O
H
97.
H
The Lewis structures for the various species are:
CO (10 e−):
CO2 (16 e−):
C
O
O
Triple bond between C and O.
O
C
CO32 (24e -):
Double bond between C and O.
2-
O
C
O
2-
O
C
O
O
2-
O
C
O
O
O
Average of 1 1/3 bond between C and O in CO32−.
H
CH3OH (14 e−):
H
C
H
O
Single bond between C and O.
H
As the number of bonds increases between two atoms, bond strength increases, and bond
length decreases. With this in mind, then:
Longest → shortest C – O bond: CH3OH > CO32− > CO2 > CO
Weakest → strongest C – O bond: CH3OH < CO32− < CO2 < CO
98.
H2NOH (14 e−)
H
N
H
O
H
Single bond between N and O
CHAPTER 8
BONDING: GENERAL CONCEPTS
N 2O (16 e-):
N
N
N
O
N
275
O
N
N
O
Average of a double bond between N and O
N O+ (10 e-):
N
N O2- (18 e-):
O N
+
O
Triple bond between N and O
-
O
O N
O
-
Average of 1 1/2 bond between N and O
N O3- (24 e-):
-
O
O
O
-
O
N
N
O
-
O
N
O
O
O
Average of 1 1/3 bond between N and O
From the Lewis structures, the order from shortest → longest N‒O bond is:
NO+ < N2O < NO2− < NO3− < H2NOH
Formal Charge
99.
BF3 has 3 + 3(7) = 24 valence electrons. The two Lewis structures to consider are:
0
F
F
+1
B
-1
F
F
B
0
0
F
0
0
F
0
The formal charges for the various atoms are assigned in the Lewis structures. Formal charge
= number of valence electrons on free atom − number of lone pair electrons on atoms −
1/2 (number of shared electrons of atom). For B in the first Lewis structure, formal charge
(FC) = 3 − 0 − 1/2(8) = −1. For F in the first structure with the double bond, FC = 7 − 4 −
1/2(4) = +1. The others all have a formal charge equal to zero [FC = 7 – 6 – 1/2(2) = 0].
The first Lewis structure obeys the octet rule but has a +1 formal charge on the most
electronegative element there is, fluorine, and a negative formal charge on a much less
electronegative element, boron. This is just the opposite of what we expect: negative formal
charge on F and positive formal charge on B. The other Lewis structure does not obey the
octet rule for B but has a zero formal charge on each element in BF3. Because structures
generally want to minimize formal charge, then BF3 with only single bonds is best from a
formal charge point of view.
276
100.
CHAPTER 8
C
BONDING: GENERAL CONCEPTS
Carbon: FC = 4 − 2 − 1/2(6) = −1; oxygen: FC = 6 − 2 − 1/2(6) = +1
O
Electronegativity predicts the opposite polarization. The two opposing effects seem to
partially cancel to give a much less polar molecule than expected.
101.
See Exercise 84 for the Lewis structures of POCl3, SO42−, ClO4− and PO43−. All these
compounds/ions have similar Lewis structures to those of SO2Cl2 and XeO4 shown below.
Formal charge = [number of valence electrons on free atom] − [number of lone pair electrons
on atom + 1/2(number of shared electrons of atom)].
a. POCl3: P, FC = 5 − 1/2(8) = +1
b. SO42−: S, FC = 6 − 1/2(8) = +2
c. ClO4−: Cl, FC = 7 − 1/2(8) = +3
d. PO43−: P, FC = 5 − 1/2(8) = +1
e. SO2Cl2, 6 + 2(6) + 2(7) = 32 e−
f.
XeO4, 8 + 4(6) = 32 eO
O
Cl
S
O
Cl
Xe
O
O
O
S, FC = 6 – 1/2(8) = +2
g. ClO3−, 7 + 3(6) + 1 = 26 e−
Xe, FC = 8 – 1/2(8) = +4
h. NO43−, 5 + 4(6) + 3 = 32 e−
3-
O
O Cl
O
O N
O
O
Cl, FC = 7 – 2 – 1/2(6) = +2
102.
O
N, FC = 5 – 1/2(8) = +1
For SO42−, ClO4−, PO43− and ClO3−, only one of the possible resonance structures is drawn.
a. Must have five bonds to P to minimize
formal charge of P. The best choice is
to form a double bond to O since this
will give O a formal charge of zero,
and single bonds to Cl for the same
reason.
O
O
Cl
P
Cl
b. Must form six bonds to S to minimize
formal charge of S.
Cl
P, FC = 0
O S O
O
2-
S, FC = 0
CHAPTER 8
BONDING: GENERAL CONCEPTS
c. Must form seven bonds to Cl
to minimize formal charge.
277
d. Must form five bonds to P to
to minimize formal charge.
O
O
Cl
O
O
P
O
P, FC = 0
O
O
e.
f.
O
Cl
3-
O
Cl, FC = O
S, FC = 0
Cl, FC = 0
O, FC = 0
Cl
S
O
O
O
Xe
O
Xe, FC = 0
O
g.
O
O Cl
Cl, FC = 0
O
h. We can’t. The following structure has a zero formal charge for N:
3-
O
O N
O
O
but N does not expand its octet. We wouldn’t expect this resonance form to exist.
103.
O2F2 has 2(6) + 2(7) = 26 valence e−. The formal charge and oxidation number (state) of
each atom is below the Lewis structure of O2F2.
F
O
O
F
Formal Charge
0
0
0
0
Oxid. Number
-1
+1
+1
-1
Oxidation states are more useful when accounting for the reactivity of O2F2. We are forced to
assign +1 as the oxidation state for oxygen due to the bonding to fluorine. Oxygen is very
electronegative, and +1 is not a stable oxidation state for this element.
278
104.
CHAPTER 8
BONDING: GENERAL CONCEPTS
OCN− has 6 + 4 + 5 + 1 = 16 valence electrons.
Formal
charge
O
C
N
O
C
N
O
C
N
0
0
-1
-1
0
0
+1
0
-2
Only the first two resonance structures should be important. The third places a positive
formal charge on the most electronegative atom in the ion and a –2 formal charge on N.
CNO− will also have 16 valence electrons.
Formal
charge
C
N
O
C
N
O
C
N
O
-2
+1
0
-1
+1
-1
-3
+1
+1
All the resonance structures for fulminate (CNO−) involve greater formal charges than in
cyanate (OCN−), making fulminate more reactive (less stable).
105.
SCl, 6 + 7 = 13; the formula could be SCl (13 valence electrons), S2Cl2 (26 valence
electrons), S3Cl3 (39 valence electrons), etc. For a formal charge of zero on S, we will need
each sulfur in the Lewis structure to have two bonds to it and two lone pairs [FC = 6 – 4 –
1/2(4) = 0]. Cl will need one bond and three lone pairs for a formal charge of zero [FC = 7 –
6 – 1/2(2) = 0]. Since chlorine wants only one bond to it, it will not be a central atom here.
With this in mind, only S2Cl2 can have a Lewis structure with a formal charge of zero on all
atoms. The structure is:
S
Cl
106.
S
Cl
The nitrogen-nitrogen bond length of 112 pm is between a double (120 pm) and a triple (110
pm) bond. The nitrogen-oxygen bond length of 119 pm is between a single (147 pm) and a
double bond (115 pm). The third resonance structure shown below doesn’t appear to be as
important as the other two since there is no evidence from bond lengths for a nitrogen-oxygen
triple bond or a nitrogen-nitrogen single bond as in the third resonance form. We can
adequately describe the structure of N2O using the resonance forms:
N
N
O
N
N
O
Assigning formal charges for all three resonance forms:
N
N
O
N
N
O
N
N
O
-1
+1
0
0
+1
-1
-2
+1
+1
CHAPTER 8
BONDING: GENERAL CONCEPTS
279
For:
, FC = 5 - 4 - 1/2(4) = -1
N
N
, FC = 5 - 1/2(8) = +1 , Same for
, FC = 5 - 6 - 1/2(2) = -2 ;
N
O ,
FC = 6 - 4 - 1/2(4) = 0 ;
O ,
FC = 6 - 2 - 1/2(6) = +1
N
O
N
and
,
FC = 5 - 2 - 1/2(6) = 0
,
FC = 6 - 6 - 1/2(2) = -1
N
We should eliminate N‒N≡O because it has a formal charge of +1 on the most
electronegative element (O). This is consistent with the observation that the N‒N bond is
between a double and triple bond and that the N‒O bond is between a single and double
bond.
107.
For formal charge values of zero:
(1) each carbon in the structure has 4 bonding pairs of electrons and no lone pairs;
(2) each N has 3 bonding pairs of electrons and 1 lone pair of electrons;
(3) each O has 2 bonding pairs of electrons and 2 lone pairs of electrons;
(4) each H is attached by only a single bond (1 bonding pair of electrons).
Following these guidelines, the Lewis structure is:
H
H
C
N
C
C
C
O
108.
H
O
C
H
H
For a formal charge of zero, carbon atoms in the structure will all satisfy the octet rule by
forming four bonds (with no lone pairs). Oxygen atoms have a formal charge of zero by
forming two bonds and having two lone pairs of electrons. Hydrogen atoms have a formal
charge of zero by forming a single bond (with no lone pairs). Following these guidelines,
two resonance structures can be drawn for benzoic acid (see next page).
280
CHAPTER 8
BONDING: GENERAL CONCEPTS
O
O
C
H
H
H
C
C
C
C
C
C
H
O
C
H
H
C
C
C
C
H
H
O
H
H
C
H
C
H
Molecular Structure and Polarity
109.
The first step always is to draw a valid Lewis structure when predicting molecular structure.
When resonance is possible, only one of the possible resonance structures is necessary to
predict the correct structure because all resonance structures give the same structure. The
Lewis structures are in Exercises 83 and 89. The structures and bond angles for each follow.
83:
a. CCl4: tetrahedral, 109.5°
b. NCl3: trigonal pyramid, <109.5°
c. SeCl2: V-shaped or bent, <109.5°
d. ICl:
linear, but there is no bond
angle present
Note: NCl3 and SeCl2 both have lone pairs of electrons on the central atom that result in bond
angles that are something less than predicted from a tetrahedral arrangement (109.5°).
However, we cannot predict the exact number. For the solutions manual, we will insert a less
than sign to indicate this phenomenon. For bond angles equal to 120°, the lone pair
phenomenon isn’t as significant as compared to smaller bond angles. For these molecules,
for example, NO2−, we will insert an approximate sign in front of the 120° to note that there
may be a slight distortion from the VSEPR predicted bond angle.
89:
a. NO2−: V-shaped, ≈120°; NO3−: trigonal planar, 120°
N2O4: trigonal planar, 120° about both N atoms
b. OCN−, SCN−, and N3− are all linear with 180° bond angles.
110.
See Exercises 84 and 90 for the Lewis structures.
84:
a. All are tetrahedral; 109.5°
b. All are trigonal pyramid; <109.5°
c. All are V-shaped; <109.5°
90:
O3 and SO2 are V-shaped (or bent) with a bond angle ≈120°. SO3 is trigonal planar
with 120° bond angles.
CHAPTER 8
111.
BONDING: GENERAL CONCEPTS
281
From the Lewis structures (see Exercise 87), Br3− would have a linear molecular structure,
ClF3 would have a T-shaped molecular structure, and SF4 would have a see-saw molecular
structure. For example, consider ClF3 (28 valence electrons):
The central Cl atom is surrounded by five electron pairs, which
requires a trigonal bipyramid geometry. Since there are three bonded
atoms and two lone pairs of electrons about Cl, we describe the
molecular structure of ClF3 as T-shaped with predicted bond angles
of about 90°. The actual bond angles will be slightly less than 90°
due to the stronger repulsive effect of the lone-pair electrons as
compared to the bonding electrons.
F
Cl
F
F
112.
From the Lewis structures (see Exercise 88), XeF4 would have a square planar molecular
structure, and ClF5 would have a square pyramid molecular structure.
113.
a. SeO3, 6 + 3(6) = 24 eO
120
o
120
Se
O
O
Se
Se
O
O
120
O
o
O
O
O
o
SeO3 has a trigonal planar molecular structure with all bond angles equal to 120°. Note
that any one of the resonance structures could be used to predict molecular structure and
bond angles.
b. SeO2, 6 + 2(6) = 18 eSe
Se
O
O
≈ 120
O
O
o
SeO2 has a V-shaped molecular structure. We would expect the bond angle to be
approximately 120° as expected for trigonal planar geometry.
Note: Both SeO3 and SeO2 structures have three effective pairs of electrons about the central
atom. All of the structures are based on a trigonal planar geometry, but only SeO3 is described as having a trigonal planar structure. Molecular structure always describes the relative
positions of the atoms.
114.
a. PCl3 has 5 + 3(7) = 26 valence electrons.
b. SCl2 has 6 + 2(7) = 20 valence
electrons.
S
P
Cl
Cl
Cl
Cl
Cl
Trigonal pyramid; all angles are <109.5°.
V-shaped; angle is <109.5°.
282
CHAPTER 8
BONDING: GENERAL CONCEPTS
c. SiF4 has 4 + 4(7) = 32 valence electrons.
F
Tetrahedral; all angles are 109.5°.
Si
F
F
F
Note: In PCl3, SCl2, and SiF4, there are four pairs of electrons about the central atom in each
case in this exercise. All of the structures are based on a tetrahedral geometry, but only SiF4
has a tetrahedral structure. We consider only the relative positions of the atoms when
describing the molecular structure.
115.
a. XeCl2 has 8 + 2(7) = 22 valence electrons.
Cl
Xe
Cl
180o
There are five pairs of electrons about the central Xe atom. The structure will be based
on a trigonal bipyramid geometry. The most stable arrangement of the atoms in XeCl2 is
a linear molecular structure with a 180° bond angle.
b. ICl3 has 7 + 3(7) = 28 valence electrons.
Cl
≈ 90o
I
Cl
Cl
T-shaped; the ClICl angles are ≈90°. Since the lone pairs will take
up more space, the ClICl bond angles will probably be slightly less
than 90°.
≈ 90o
c. TeF4 has 6 + 4(7) = 34
valence electrons.
d. PCl5 has 5 + 5(7) = 40
valence electrons.
90o
F
Cl
F
≈ 120o
Te
F
F
o
120
Cl
P
Cl
Cl
Cl
≈ 90o
See-saw or teeter-totter
or distorted tetrahedron
Trigonal bipyramid
All the species in this exercise have five pairs of electrons around the central atom. All
the structures are based on a trigonal bipyramid geometry, but only in PCl5 are all the
CHAPTER 8
BONDING: GENERAL CONCEPTS
283
pairs, bonding pairs. Thus PCl5 is the only one for which we describe the molecular
structure as trigonal bipyramid. Still, we had to begin with the trigonal bipyramid
geometry to get to the structures (and bond angles) of the others.
116.
a. ICl5 , 7 + 5(7) = 42 e-
b. XeCl4 , 8 + 4(7) = 36 e-
≈ 90o
90o
Cl
Cl
Cl
Cl
I
Cl
Cl
Xe
90o
Cl
Cl
90o
Cl
90o
Square pyramid, ≈90° bond angles
Square planar, 90° bond angles
c. SeCl6 has 6 + 6(7) = 48 valence electrons.
Cl
Cl
Cl
Octahedral, 90° bond angles
Se
Cl
Cl
Cl
Note: All these species have six pairs of electrons around the central atom. All three
structures are based on the octahedron, but only SeCl6 has an octahedral molecular structure.
117.
SeO3 and SeO2 both have polar bonds, but only SeO2 has a dipole moment. The three bond
dipoles from the three polar Se‒O bonds in SeO3 will all cancel when summed together.
Hence SeO3 is nonpolar since the overall molecule has no resulting dipole moment. In SeO2,
the two Se‒O bond dipoles do not cancel when summed together; hence SeO2 has a net
dipole moment (is polar). Since O is more electronegative than Se, the negative end of the
dipole moment is between the two O atoms, and the positive end is around the Se atom. The
arrow in the following illustration represents the overall dipole moment in SeO2. Note that to
predict polarity for SeO2, either of the two resonance structures can be used.
Se
O
118.
O
All have polar bonds; in SiF4, the individual bond dipoles cancel when summed together, and
in PCl3 and SCl2, the individual bond dipoles do not cancel. Therefore, SiF4 has no net dipole
moment (is nonpolar), and PCl3 and SCl2 have net dipole moments (are polar). For PCl3, the
negative end of the dipole moment is between the more electronegative chlorine atoms, and
the positive end is around P. For SCl2, the negative end is between the more electronegative
Cl atoms, and the positive end of the dipole moment is around S.
284
119.
CHAPTER 8
BONDING: GENERAL CONCEPTS
All have polar bonds, but only TeF4 and ICl3 have dipole moments. The bond dipoles from
the five P‒Cl bonds in PCl5 cancel each other when summed together, so PCl5 has no net
dipole moment. The bond dipoles in XeCl2 also cancel:
Xe
Cl
Cl
Because the bond dipoles from the two Xe‒Cl bonds are equal in magnitude but point in
opposite directions, they cancel each other, and XeCl2 has no net dipole moment (is
nonpolar). For TeF4 and ICl3, the arrangement of these molecules is such that the individual
bond dipoles do not all cancel, so each has an overall net dipole moment (is polar).
120.
All have polar bonds, but only ICl5 has an overall net dipole moment. The six bond dipoles in
SeCl6 all cancel each other, so SeCl6 has no net dipole moment. The same is true for XeCl4:
Cl
Cl
Xe
Cl
Cl
When the four bond dipoles are added together, they all cancel each other, resulting in XeCl4
having no overall dipole moment (is nonpolar). ICl5 has a structure in which the individual
bond dipoles do not all cancel, hence ICl5 has a dipole moment (is polar)
121.
Molecules that have an overall dipole moment are called polar molecules, and molecules that
do not have an overall dipole moment are called nonpolar molecules.
a. OCl2, 6 + 2(7) = 20 e−
KrF2, 8 + 2(7) = 22 e−
O
Cl
O
Cl
Cl
F
Kr
F
Cl
V-shaped, polar; OCl2 is polar because
the two O−Cl bond dipoles don’t cancel
each other. The resulting dipole moment
is shown in the drawing.
Linear, nonpolar; the molecule is
nonpolar because the two Kr‒F
bond dipoles cancel each other.
CHAPTER 8
BONDING: GENERAL CONCEPTS
BeH2, 2 + 2(1) = 4 e−
H
Be
285
SO2, 6 + 2(6) = 18 e−
S
H
O
Linear, nonpolar; Be−H bond dipoles
are equal and point in opposite directions.
They cancel each other. BeH2 is nonpolar.
O
V-shaped, polar; the S−O bond dipoles
do not cancel, so SO2 is polar (has a net
dipole moment). Only one resonance
structure is shown.
Note: All four species contain three atoms. They have different structures because the
number of lone pairs of electrons around the central atom are different in each case.
b. SO3, 6 + 3(6) = 24 e−
NF3, 5 + 3(7) = 26 e-
O
N
F
S
O
F
F
O
Trigonal planar, nonpolar;
bond dipoles cancel. Only one
resonance structure is shown.
Trigonal pyramid, polar;
bond dipoles do not cancel.
IF3 has 7 + 3(7) = 28 valence electrons.
F
I
F
T-shaped, polar; bond dipoles do not cancel.
F
Note: Each molecule has the same number of atoms but different structures because of
differing numbers of lone pairs around each central atom.
c. CF4, 4 + 4(7) = 32 e−
F
F
F
Se
C
F
SeF4, 6 + 4(7) = 34 e−
F
F
Tetrahedral, nonpolar;
bond dipoles cancel.
F
F
See-saw, polar;
bond dipoles do not cancel.
286
CHAPTER 8
BONDING: GENERAL CONCEPTS
KrF4, 8 + 4(7) = 36 valence electrons
F
F
Square planar, nonpolar;
bond dipoles cancel.
Kr
F
F
Note: Again, each molecule has the same number of atoms but different structures
because of differing numbers of lone pairs around the central atom.
d. IF5, 7 + 5(7) = 42 e−
AsF5, 5 + 5(7) = 40 e−
F
F
F
F
F
F
F
As
I
F
F
F
Square pyramid, polar;
bond dipoles do not cancel.
Trigonal bipyramid, nonpolar;
bond dipoles cancel.
Note: Yet again, the molecules have the same number of atoms but different structures
because of the presence of differing numbers of lone pairs.
122.
a.
b.
O
C
S
N
C
O
H
Polar; the bond dipoles do
not cancel.
c.
Polar; the C‒O bond is a more polar
bond than the C‒S bond, so the two
bond dipoles do not cancel each other.
d.
F
F
Xe
F
C
F
Cl
Cl
Nonpolar; the two Xe‒F bond
dipoles cancel each other.
Polar; all the bond dipoles are not
equivalent, and they don’t cancel each
other.
CHAPTER 8
BONDING: GENERAL CONCEPTS
e.
287
f.
F
F
H
F
C
Se
F
F
O
H
F
Nonpolar; the six Se−F bond
dipoles cancel each other.
123.
Polar; the bond dipoles are not
equivalent, and they don’t cancel
EO3− is the formula of the ion. The Lewis structure has 26 valence electrons. Let x =
number of valence electrons of element E.
26 = x + 3(6) + 1, x = 7 valence electrons
Element E is a halogen because halogens have seven valence electrons. Some possible
identities are F, Cl, Br, and I. The EO3− ion has a trigonal pyramid molecular structure with
bond angles of less than 109.5° (<109.5°).
124.
The formula is EF2O2−, and the Lewis structure has 28 valence electrons.
28 = x + 2(7) + 6 + 2, x = 6 valence electrons for element E
Element E must belong to the Group 6A elements since E has six valence electrons. E must
also be a row 3 or heavier element since this ion has more than eight electrons around the
central E atom (row 2 elements never have more than eight electrons around them). Some
possible identities for E are S, Se, and Te. The ion has a T-shaped molecular structure with
bond angles of ≈90°.
125.
All these molecules have polar bonds that are symmetrically arranged about the central
atoms. In each molecule, the individual bond dipoles cancel each other out to give no net
overall dipole moment. All these molecules are nonpolar even though they all contain polar
bonds.
126.
XeF 2Cl2, 8 + 2(7) + 2(7) = 36 e−
Cl
F
Cl
F
F
Xe
Xe
Cl
polar
F
Cl
nonpolar
The two possible structures for XeF2Cl2 are above. In the first structure, the F atoms are 90°
apart from each other, and the Cl atoms are also 90° apart. The individual bond dipoles would
not cancel in this molecule, so this molecule is polar. In the second possible structure, the F
atoms are 180° apart, as are the Cl atoms. Here, the bond dipoles are symmetrically arranged
so they do cancel each other out, and this molecule is nonpolar. Therefore, measurement of
the dipole moment would differentiate between the two compounds. These are different
compounds and not resonance structures.
288
CHAPTER 8
BONDING: GENERAL CONCEPTS
Additional Exercises
127.
a. Radius: N+ < N < N− ; IE: N− < N < N+
N+ has the fewest electrons held by the seven protons in the nucleus, whereas N− has the
most electrons held by the seven protons. The seven protons in the nucleus will hold the
electrons most tightly in N+ and least tightly in N− . Therefore, N+ has the smallest radius
with the largest ionization energy (IE), and N− is the largest species with the smallest IE.
b. Radius: Cl+ < Cl < Se < Se− ; IE: Se− < Se < Cl < Cl+
The general trends tell us that Cl has a smaller radius than Se and a larger IE than Se.
Cl+, with fewer electron-electron repulsions than Cl, will be smaller than Cl and have a
larger IE. Se−, with more electron-electron repulsions than Se, will be larger than Se and
have a smaller IE.
c. Radius: Sr2+ < Rb+ < Br− ; IE: Br- < Rb+ < Sr2+
These ions are isoelectronic. The species with the most protons (Sr2+) will hold the
electrons most tightly and will have the smallest radius and largest IE. The ion with the
fewest protons (Br−) will hold the electrons least tightly and will have the largest radius
and smallest IE.
128.
129.
a. Na+(g) + Cl−(g) → NaCl(s)
b. NH4+(g) + Br-(g) → NH4Br(s)
c. Mg2+(g) + S2−(g) → MgS(s)
d. O2(g) → 2 O(g)
a.
HF(g) → H(g) + F(g)
H(g) → H+(g) + e−
F(g) + e− → F−(g)
HF(g) → H+(g) + F−(g)
b.
HCl(g) → H(g) + Cl(g)
H(g) → H+(g) + e−
Cl(g) + e−→ Cl−(g)
HCl(g) → H+(g) + Cl−(g)
c.
HI(g) → H(g) + I(g)
H(g) → H+(g) + e−
I(g) + e− → I−(g)
HI(g) → H+(g) + I−(g)
d.
H2O(g) → OH(g) + H(g)
H(g) → H+(g) + e−
OH(g) + e− → OH−(g)
H2O(g) → H+(g) + OH−(g)
ΔH = 565 kJ
ΔH = 1312 kJ
ΔH = −327.8 kJ
ΔH = 1549 kJ
ΔH = 427 kJ
ΔH = 1312 kJ
ΔH = −348.7 kJ
ΔH = 1390. kJ
ΔH = 295 kJ
ΔH = 1312 kJ
ΔH = −295.2 kJ
ΔH = 1312 kJ
ΔH = 467 kJ
ΔH = 1312 kJ
ΔH = −180. kJ
ΔH = 1599 kJ
CHAPTER 8
130.
BONDING: GENERAL CONCEPTS
289
CO32− has 4 + 3(6) + 2 = 24 valence electrons.
2-
O
O
O
2-
O
C
C
O
2-
O
C
O
O
O
HCO3− has 1 + 4 + 3(6) + 1 = 24 valence electrons.
H
-
O
H
C
C
O
-
O
O
O
O
H2CO3 has 2(1) + 4 + 3(6) = 24 valence electrons.
O
C
O
O
H
H
The Lewis structures for the reactants and products are:
O
C
O
H
H
O
H
+
O
C
O
O
H
Bonds broken:
2 C‒O (358 kJ/mol)
1 O‒H (467 kJ/mol)
Bonds formed:
1 C=O (799 kJ/mol)
1 O‒H (467 kJ/mol)
ΔH = 2(358) + 467 − [799 + 467] = −83 kJ; the carbon-oxygen double bond is stronger than
two carbon-oxygen single bonds; hence CO2 and H2O are more stable than H2CO3.
131.
The stable species are:
a. NaBr: In NaBr2, the sodium ion would have a 2+ charge, assuming that each bromine
has a 1− charge. Sodium doesn’t form stable Na2+ ionic compounds.
b. ClO4−: ClO4 has 31 valence electrons, so it is impossible to satisfy the octet rule for all
atoms in ClO4. The extra electron from the 1− charge in ClO4- allows for complete octets
for all atoms.
290
CHAPTER 8
BONDING: GENERAL CONCEPTS
c. XeO4: We can’t draw a Lewis structure that obeys the octet rule for SO4 (30 electrons),
unlike XeO4 (32 electrons).
d. SeF4: Both compounds require the central atom to expand its octet. O is too small and
doesn’t have low-energy d orbitals to expand its octet (which is true for all row 2
elements).
132.
a. All have 24 valence electrons and the same number of atoms in the formula. All have
the same resonance Lewis structures; the structures are all trigonal planar with 120° bond
angles. The Lewis structures for NO3− and CO32− will be the same as the three SO3 Lewis
structures shown below.
O
O
S
S
O
O
O
O
S
O
O
O
b. All have 18 valence electrons and the same number of atoms. All have the same
resonance Lewis structures; the molecular structures are all V-shaped with ≈120° bond
angles. O3 and SO2 have the same two Lewis structures as is shown for NO2−.
N
N
O
133.
O
O
a. XeCl4, 8 + 4(7) = 36 e−
Cl
XeCl2, 8 + 2(7) = 22 e−
Cl
Xe
Cl
O
Cl
Square planar, 90°, nonpolar
Cl
Xe
Cl
Linear, 180°, nonpolar
Both compounds have a central Xe atom with lone pairs and terminal Cl atoms, and both
compounds do not satisfy the octet rule. In addition, both are nonpolar because the
Xe−Cl bond dipoles and lone pairs around Xe are arranged in such a manner that they
cancel each other out. The last item in common is that both have 180° bond angles.
Although we haven’t emphasized this, the bond angles between the Cl atoms on the
diagonal in XeCl4 are 180° apart from each other.
b. All of these are polar covalent compounds. The bond dipoles do not cancel out each other
when summed together. The reason the bond dipoles are not symmetrically arranged in
these compounds is that they all have at least one lone pair of electrons on the central
atom, which disrupts the symmetry. Note that there are molecules that have lone pairs
and are nonpolar, e.g., XeCl4 and XeCl2 in the preceding problem. A lone pair on a
central atom does not guarantee a polar molecule.
CHAPTER 8
134.
BONDING: GENERAL CONCEPTS
291
The general structure of the trihalide ions is:
X
X
X
Bromine and iodine are large enough and have low-energy, empty d orbitals to accommodate
the expanded octet. Fluorine is small, and its valence shell contains only 2s and 2p orbitals
(four orbitals) and cannot expand its octet. The lowest-energy d orbitals in F are 3d; they are
too high in energy compared with 2s and 2p to be used in bonding.
135.
Yes, each structure has the same number of effective pairs around the central atom, giving the
same predicted molecular structure for each compound/ion. (A multiple bond is counted as a
single group of electrons.)
136.
a.
The C‒H bonds are assumed nonpolar since the
electronegativities of C and H are about equal.
H
C
H
Cl
Cl
δ+ δ−
C‒Cl is the charge distribution for each C‒Cl bond.
In CH2Cl2, the two individual C‒Cl bond dipoles
add together to give an overall dipole moment for
the molecule. The overall dipole will point from C
(positive end) to the midpoint of the two Cl atoms
(negative end).
In CHCl3, the C‒H bond is essentially nonpolar. The three C‒Cl bond dipoles in CHCl3
add together to give an overall dipole moment for the molecule. The overall dipole will
have the negative end at the midpoint of the three chlorines and the positive end around
the carbon.
H
C
Cl
Cl
Cl
CCl4 is nonpolar. CCl4 is a tetrahedral molecule where all four C‒Cl bond dipoles cancel
when added together. Let’s consider just the C and two of the Cl atoms. There will be a
net dipole pointing in the direction of the middle of the two Cl atoms.
C
Cl
Cl
There will be an equal and opposite dipole arising from the other two Cl atoms.
292
CHAPTER 8
BONDING: GENERAL CONCEPTS
Combining:
Cl
Cl
C
Cl
Cl
The two dipoles cancel, and CCl4 is nonpolar.
b. CO2 is nonpolar. CO2 is a linear molecule with two equivalence bond dipoles that cancel.
N2O, which is also a linear molecule, is polar because the nonequivalent bond dipoles do
not cancel.
N
δ+
N
δ−
O
c. NH3 is polar. The 3 N‒H bond dipoles add together to give a net dipole in the direction of
the lone pair. We would predict PH3 to be nonpolar on the basis of electronegativitity,
i.e., P‒H bonds are nonpolar. However, the presence of the lone pair makes the PH3
molecule slightly polar. The net dipole is in the direction of the lone pair and has a
magnitude about one third that of the NH3 dipole.
δ−
N
δ+
N‒H
137.
H
H
P
H
H
H
H
TeF5− has 6 + 5(7) + 1 = 42 valence electrons.
-
F
F
F
Te
F
F
The lone pair of electrons around Te exerts a stronger repulsion than the bonding pairs of
electrons. This pushes the four square-planar F atoms away from the lone pair and reduces
the bond angles between the axial F atom and the square-planar F atoms.
138.
C ≡ O (1072 kJ/mol); N ≡ N (941 kJ/mol); CO is polar, whereas N2 is nonpolar. This may
lead to a great reactivity for the CO bond.
CHAPTER 8
BONDING: GENERAL CONCEPTS
293
ChemWork Problems
The answers to the problems 139-146 (or a variation to these problems) are found in OWL. These
problems are also assignable in OWL.
Challenge Problems
147.
a. There are two attractions of the form
(+1)(−1)
, where r = 1 × 10−10 m = 0.1 nm.
r
 (+1)(−1) 
−18
−18
V = 2 × (2.31 × 10−19 J nm) 
 = −4.62 × 10 J = −5 × 10 J
0
.
1
nm


b. There are four attractions of +1 and −1 charges at a distance of 0.1 nm from each other.
The two negative charges and the two positive charges repel each other across the
diagonal of the square. This is at a distance of 2 × 0.1 nm.
 (+1)(−1) 
−19
V = 4 × (2.31 × 10−19) 
 + 2.31 × 10
 0.1 
 (+1)(+1) 


 2 (0.1) 
 (−1)(−1) 
+ 2.31 × 10−19 

 2 (0.1) 
V = −9.24 × 10−18 J + 1.63 × 10−18 J + 1.63 × 10−18 J = −5.98 × 10−18 J = −6 × 10−18 J
Note: There is a greater net attraction in arrangement b than in a.
(IE − EA)
148.
F
Cl
Br
I
2006 kJ/mol
1604 kJ/mol
1463 kJ/mol
1302 kJ/mol
(IE − EA)/502
4.0
3.2
2.9
2.6
EN (text)
2006/502 = 4.0
4.0
3.0
2.8
2.5
The values calculated from ionization energies and electron affinities show the same trend as
(and agree fairly closely) with the values given in the text.
149.
The reaction is:
1/2 I2(g) + 1/2 Cl2(g) → ICl(g)
ΔH of = ?
Using Hess’s law:
1/2 I2(s) → 1/2 I2(g)
1/2 I2(g) → I (g)
1/2 Cl2(g) → Cl(g)
I(g) + Cl(g) → ICl(g)
1/2 I2(s) + 1/2 Cl2(g) → ICl(g)
∆H = 1/2(62 kJ)
∆H = 1/2(149 kJ)
∆H = 1/2(239 kJ)
∆H = −208 kJ
(Appendix 4)
(Table 8.4)
(Table 8.4)
(Table 8.4)
∆H = 17 kJ so ΔH of = 17 kJ/mol
294
150.
CHAPTER 8
2 Li+(g) + 2 Cl−(g) →
2 Li(g) →
2 Li(s) →
2 HCl(g) →
2 Cl(g) + 2 e− →
2 H(g) →
2 LiCl(s)
2 Li+(g) + 2 e−
2 Li(g)
2 H(g) + 2 Cl(g)
2 Cl−(g)
H2(g)
2 Li(s) + 2 HCl(g) → 2 LiCl(s) + H2(g)
BONDING: GENERAL CONCEPTS
∆H = 2(−829 kJ)
∆H = 2(520. kJ)
∆H = 2(166 kJ)
∆H = 2(427 kJ)
∆H = 2(−349 kJ)
∆H = −(432 kJ)
∆H = −562 kJ
151.
See Figure 8.11 to see the data supporting MgO as an ionic compound. Note that the lattice
energy is large enough to overcome all of the other processes (removing two electrons from
Mg, etc.). The bond energy for O2 (247 kJ/mol) and electron affinity (737 kJ/mol) are the
same when making CO. However, ionizing carbon to form a C2+ ion must be too large. See
Figure 7.32 to see that the first ionization energy for carbon is about 350 kJ/mol greater than
the first ionization energy for magnesium. If all other numbers were equal, the overall energy
change would be down to ~250 kJ/mol (see Figure 8.11). It is not unreasonable to assume
that the second ionization energy for carbon is more than 250 kJ/mol greater than the second
ionization energy of magnesium. This would result in a positive ∆H value for the formation
of CO as an ionic compound. One wouldn’t expect CO to be ionic if the energetics were
unfavorable.
152.
a. (1) Removing an electron from the metal: ionization energy, positive (∆H > 0)
(2) Adding an electron to the nonmetal: electron affinity, often negative (∆H < 0)
(3) Allowing the metal cation and nonmetal anion to come together: lattice energy,
negative (∆H < 0)
b. Often the sign of the sum of the first two processes is positive (or unfavorable). This is
especially true due to the fact that we must also vaporize the metal and often break a
bond on a diatomic gas. For example, the ionization energy for Na is +495 kJ/mol, and
the electron affinity for F is −328 kJ/mol. Overall, the energy change is +167 kJ/mol
(unfavorable).
c. For an ionic compound to form, the sum must be negative (exothermic).
d. The lattice energy must be favorable enough to overcome the endothermic process of
forming the ions; i.e., the lattice energy must be a large negative quantity.
e. While Na2Cl (or NaCl2) would have a greater lattice energy than NaCl, the energy to
make a Cl2− ion (or Na2+ ion) must be larger (more unfavorable) than what would be
gained by the larger lattice energy. The same argument can be made for MgO compared
to MgO2 or Mg2O. The energy to make the ions is too unfavorable or the lattice energy is
not favorable enough, and the compounds do not form.
153.
As the halogen atoms get larger, it becomes more difficult to fit three halogen atoms around
the small nitrogen atom, and the NX3 molecule becomes less stable.
CHAPTER 8
154.
BONDING: GENERAL CONCEPTS
295
a. I.
H
H * O
H
C
H
+
C
* H
O
C
N
C
C
H
H
H
*
H * C
N
Bonds broken (*):
H
Bonds formed (*):
1 C‒O (358 kJ)
1 C‒H (413 kJ)
1 O‒H (467 kJ)
1 C‒C (347 kJ)
ΔHI = 358 kJ + 413 kJ − (467 kJ + 347 kJ) = −43 kJ
II.
OH H
H
*
C * C
*
H
H
H
C
N
H
C
*
H
Bonds broken (*):
+
C
C
H * O
H
N
Bonds formed (*):
1 C‒O (358 kJ/mol)
1 C‒H (413 kJ/mol)
1 C‒C (347 kJ/mol)
1 H‒O (467 kJ/mol)
1 C=C (614 kJ/mol)
ΔHII = 358 kJ + 413 kJ + 347 kJ − [467 kJ + 614 kJ] = +37 kJ
ΔHoverall = ΔHI + ΔHII = −43 kJ + 37 kJ = −6 kJ
b.
H
H
4
C
H
H
H
C
C
H
+ 6 NO
H
Bonds broken:
4 × 3 C‒H (413 kJ/mol)
6 N=O (630. kJ/mol)
4
C
C
+ 6 H
C
H
N
H
Bonds formed:
4 C≡N (891 kJ/mol)
6 × 2 H‒O (467 kJ/mol)
1 N≡N (941 kJ/mol)
ΔH = 12(413) + 6(630.) − [4(891) + 12(467) + 941] = −1373 kJ
O
H + N
N
296
CHAPTER 8
BONDING: GENERAL CONCEPTS
c.
H
H
2
H
H
C
C
H
C
H
+ 2
H
N
H
H
+ 3 O2
C
C
2
Bonds broken:
+ 6
C
H
H
N
H
O
H
H
Bonds formed:
2 C≡N (891 kJ/mol)
6 × 2 O‒H (467 kJ/mol)
2 × 3 C‒H (413 kJ/mol)
2 × 3 N‒H (391 kJ/mol)
3 O=O (495 kJ/mol)
ΔH = 6(413) + 6(391) + 3(495) − [2(891) + 12(467)] = −1077 kJ
d. Because both reactions are highly exothermic, the high temperature is not needed to
provide energy. It must be necessary for some other reason. The reason is to increase
the speed of the reaction. This is discussed in Chapter 12 on kinetics.
155.
a. i.
C6H6N12O12 → 6 CO + 6 N2 + 3 H2O + 3/2 O2
The NO2 groups are assumed to have one N‒O single bond and one N=O double
bond, and each carbon atom has one C‒H single bond. We must break and form all
bonds.
Bonds broken:
Bonds formed:
3 C‒C (347 kJ/mol)
6 C‒H (413 kJ/mol)
12 C‒N (305 kJ/mol)
6 N‒N (160. kJ/mol)
6 N‒O (201 kJ/mol)
6 N=O (607 kJ/mol)
ΣDbroken = 12,987 kJ
6 C≡O (1072 kJ/mol)
6 N≡N (941 kJ/mol)
6 H‒O (467 kJ/mol)
3/2 O=O (495 kJ/mol)
ΣDformed = 15,623 kJ
ΔH = ΣDbroken − ΣDformed = 12,987 kJ − 15,623 kJ = −2636 kJ
ii. C6H6N12O12 → 3 CO + 3 CO2 + 6 N2 + 3 H2O
Note: The bonds broken will be the same for all three reactions.
Bonds formed:
3 C≡O (1072 kJ/mol)
6 C=O (799 kJ/mol)
6 N≡N (941 kJ/mol)
6 H‒O (467 kJ/mol)
ΣDformed = 16,458 kJ
ΔH = 12,987 kJ −16,458 kJ = −3471 kJ
CHAPTER 8
BONDING: GENERAL CONCEPTS
297
iii. C6H6N12O12 → 6 CO2 + 6 N2 + 3 H2
Bonds formed:
12 C=O (799 kJ/mol)
6 N≡N (941 kJ/mol)
3 H‒H (432 kJ/mol)
ΣDformed = 16,530. kJ
ΔH = 12,987 kJ − 16,530. kJ = −3543 kJ
b. Reaction iii yields the most energy per mole of CL-20, so it will yield the most energy
per kilogram.
− 3543 kJ
1 mol
1000 g
×
×
= −8085 kJ/kg
mol
438.23 g
kg
156.
We can draw resonance forms for the anion after the loss of H+, we can argue that the extra
stability of the anion causes the proton to be more readily lost, i.e., makes the compound a
better acid.
a.
-
O
H
C
-
O
O
H
C
O
b.
O
CH3
C
-
O
CH
C
CH3
O
CH3
C
O
CH
C
CH3
O
CH3
C
O
CH
C
CH3
298
CHAPTER 8
BONDING: GENERAL CONCEPTS
c.
O
O
O
O
O
In all three cases, extra resonance forms can be drawn for the anion that are not possible
when the H+ is present, which leads to enhanced stability.
157.
For carbon atoms to have a formal charge of zero, each C atom must satisfy the octet rule by
forming four bonds (with no lone pairs). For nitrogen atoms to have a formal charge of zero,
each N atom must satisfy the octet rule by forming three bonds and have one lone pair of
electrons. For oxygen atoms to have a formal charge of zero, each O atom must satisfy the
octet rule by forming two bonds and have two lone pairs of electrons. With these bonding
requirements in mind, then the Lewis structure of histidine, where all atoms have a formal
charge of zero, is:
H
2
H
C
N
C
N
H
C
H
N
C
C
H
H
O
C H
H
1
O
H
We would expect 120° bond angles about the carbon atom labeled 1 and ≈109.5° bond angles
about the nitrogen atom labeled 2. The nitrogen bond angles should be slightly smaller than
109.5° due to the lone pair of electrons on nitrogen.
158.
This molecule has 30 valence electrons. The only C–N bond that can possibly have a doublebond character is the N bound to the C with O attached. Double bonds to the other two C–N
bonds would require carbon in each case to have 10 valence electrons (which carbon never
does). The resonance structures are:
CHAPTER 8
BONDING: GENERAL CONCEPTS
O
H
C
H
N
C
299
O
H
H
C
H
N
H
H
C
H
H
H
H
H
159.
C
C
H
H
a. BrFI2, 7 +7 + 2(7) = 28 e−; two possible structures exist with Br as the central atom; each
has a T-shaped molecular structure.
I
F
Br
I
Br
F
I
I
90° bond angles between I atoms
180° bond angles between I atoms
b. XeO2F2, 8 + 2(6) + 2(7) = 34 e−; three possible structures exist with Xe as the central
atom; each has a see-saw molecular structure.
O
O
F
Xe
F
O
O
Xe
Xe
F
F
O
O
180° bond angle
between O atoms
90° bond angle
between O atoms
F
F
120° bond angle
between O atoms
c. TeF2Cl3−; 6 + 2(7) + 3(7) + 1 = 42 e−; three possible structures exist with Te as the
central atom; each has a square pyramid molecular structure.
Cl
F
Cl
Cl
F
F
Te
Te
Cl
Cl
F
One F is 180° from
the lone pair.
Cl
Cl
Cl
Te
F
Both F atoms are 90°
from the lone pair and 90°
from each other.
Cl
F
Both F atoms are 90°
from the lone pair and 180°
from each other.
300
160.
CHAPTER 8
BONDING: GENERAL CONCEPTS
The skeletal structure of caffeine is:
H
O
H
H
C
H
H
C
C
H
N
N
C
C
C
C
O
N
N
H
H
C
H
H
For a formal charge of zero on all atoms, the bonding requirements are:
(1) four bonds and no lone pairs for each carbon atom;
(2) three bonds and one lone pair for each nitrogen atom;
(3) two bonds and two lone pairs for each oxygen atom;
(4) one bond and no lone pairs for each hydrogen atom.
Following these guidelines gives a Lewis structure that has a formal charge of zero for all the
atoms in the molecule. The Lewis structure is:
H
O
H
H
C
H
H
C
H
C
N
N
C
C
C
C
N
O
H
C
H
N
H
H
CHAPTER 8
BONDING: GENERAL CONCEPTS
301
Integrative Problems
161.
Assuming 100.00 g of compound: 42.81 g F =
1 mol X
= 2.253 mol F
19.00 g F
The number of moles of X in XF5 is: 2.53 mol F ×
1 mol X
= 0.4506 mol X
5 mol F
This number of moles of X has a mass of 57.19 g (= 100.00 g – 42.81 g). The molar mass of
X is:
57.19 g X
= 126.9 g/mol; this is element I and the compound is IF5.
0.4506 mol X
IF5, 7 + 5(7) = 42 e−
F
F
F
I
F
162.
The molecular structure is square pyramid.
F
If X2− has a configuration of [Ar]4s23d104p6, then X must have a configuration with two fewer
electrons, [Ar]4s23d104p4. This is element Se.
SeCN−, 6 + 4 + 5 + 1 = 16 e−
Se
163.
Se
N
C
C
N
Se
C
N
The elements are identified by their electron configurations:
[Ar]4s13d5 = Cr; [Ne]3s23p3 = P; [Ar]4s23d104p3 = As; [Ne]3s23p5 = Cl
Following the electronegativity trend, the order is Cr < As < P < Cl.
Marathon Problem
164.
Compound A: This compound is a strong acid (part g). HNO3 is a strong acid and is
available in concentrated solutions of 16 M (part c). The highest possible oxidation state of
nitrogen is +5, and in HNO3, the oxidation state of nitrogen is +5 (part b). Therefore,
compound A is most likely HNO3. The Lewis structures for HNO3 are:
O
H
O
N
O
H
N
O
O
O
302
CHAPTER 8
BONDING: GENERAL CONCEPTS
Compound B: This compound is basic (part g) and has one nitrogen (part b). The formal
charge of zero (part b) tells us that there are three bonds to the nitrogen and that the nitrogen
has one lone pair. Assuming compound B is monobasic, then the data in part g tell us that the
molar mass of B is 33.0 g/mol (21.98 mL of 1.000 M HCl = 0.02198 mol HCl; thus there are
0.02198 mol of B; 0.726 g/0.02198 mol = 33.0 g/mol). Because this number is rather small,
it limits the possibilities. That is, there is one nitrogen, and the remainder of the atoms are O
and H. Since the molar mass of B is 33.0 g/mol, then only one O oxygen atom can be
present. The N and O atoms have a combined molar mass of 30.0 g/mol; the rest is made up
of hydrogens (3 H atoms), giving the formula NH3O. From the list of Kb values for weak
bases in Appendix 5 of the text, compound B is most likely NH2OH. The Lewis structure is:
H
N
O
H
H
Compound C: From parts a and f and assuming compound A is HNO3 , then compound C
contains the nitrate ion, NO3-. Because part b tells us that there are two nitrogens, the other
ion needs to have one N atom and some H atoms. In addition, compound C must be a weak
acid (part g), which must be due to the other ion since NO3- has no acidic properties. Also,
the nitrogen atom in the other ion must have an oxidation state of -3 (part b) and a formal
charge of +1. The ammonium ion fits the data. Thus compound C is most likely NH4NO3.
A Lewis structure is:
+
H
H
N
H
-
O
N
H
O
O
Note: Two more resonance structures can be drawn for NO3-.
Compound D: From part f, this compound has one less oxygen atom than compound C; thus
NH4NO2 is a likely formula. Data from part e confirm this. Assuming 100.0 g of compound,
we have:
43.7 g N × 1 mol/14.01 g = 3.12 mol N
50.0 g O × 1 mol/16.00 g = 3.12 mol O
6.3 g H × 1 mol/1.008 g = 6.25 mol H
There is a 1 : 1 : 2 mole ratio among N to O to H. The empirical formula is NOH2, which has
an empirical formula mass of 32.0 g/mol.
Molar mass =
2.86 g/L(0.08206 L atm/K • mol)(273 K)
dRT
=
= 64.1 g/mol
1.00 atm
P
CHAPTER 8
BONDING: GENERAL CONCEPTS
303
For a correct molar mass, the molecular formula of compound D is N2O2H4 or NH4NO2. A
Lewis structure is:
H
N
-
+
H
N
H
O
O
H
Note: One more resonance structure for NO2− can be drawn.
Compound E: A basic solution (part g) that is commercially available at 15 M (part c) is
ammonium hydroxide (NH4OH). This is also consistent with the information given in parts b
and d. The Lewis structure for NH4OH is:
+
H
H
N
H
H
O
H
CHAPTER 9
COVALENT BONDING: ORBITALS
Questions
9.
In hybrid orbital theory, some or all of the valence atomic orbitals of the central atom in a
molecule are mixed together to form hybrid orbitals; these hybrid orbitals point to where the
bonded atoms and lone pairs are oriented. The sigma bonds are formed from the hybrid
orbitals overlapping head to head with an appropriate orbital from the bonded atom. The π
bonds, in hybrid orbital theory, are formed from unhybridized p atomic orbitals. The p
orbitals overlap side to side to form the π bond, where the π electrons occupy the space above
and below a line joining the atoms (the internuclear axis). Assuming the z-axis is the
internuclear axis, then the pz atomic orbital will always be hybridized whether the
hybridization is sp, sp2, sp3, dsp3 or d2sp3. For sp hybridization, the px and py atomic orbitals
are unhybridized; they are used to form two π bonds to the bonded atom(s). For sp2
hybridization, either the px or the py atomic orbital is hybridized (along with the s and pz
orbitals); the other p orbital is used to form a π bond to a bonded atom. For sp3 hybridization,
the s and all the p orbitals are hybridized; no unhybridized p atomic orbitals are present, so no
π bonds form with sp3 hybridization. For dsp3 and d2sp3 hybridization, we just mix in one or
two d orbitals into the hybridization process. Which specific d orbitals are used is not
important to our discussion.
10.
The MO theory is a mathematical model. The allowed electron energy levels (molecular
orbitals) in a molecule are solutions to the mathematical problem. The square of the solutions
gives the shapes of the molecular orbitals. A sigma bond is an allowed energy level where the
greatest electron probability is between the nuclei forming the bond. Valence s orbitals form
sigma bonds, and if the z-axis is the internuclear axis, then valence pz orbitals also form
sigma bonds. For a molecule like HF, a sigma-bonding MO results from the combination of
the H 1s orbital and the F 2pz atomic orbital.
For π bonds, the electron density lies above and below the internuclear axis. The π bonds are
formed when px orbitals are combined (side-to-side overlap) and when py orbitals are
combined.
11.
We use d orbitals when we have to; i.e., we use d orbitals when the central atom on a
molecule has more than eight electrons around it. The d orbitals are necessary to accommodate the electrons over eight. Row 2 elements never have more than eight electrons around
them, so they never hybridize d orbitals. We rationalize this by saying there are no d orbitals
close in energy to the valence 2s and 2p orbitals (2d orbitals are forbidden energy levels).
However, for row 3 and heavier elements, there are 3d, 4d, 5d, etc. orbitals that will be close
in energy to the valence s and p orbitals. It is row 3 and heavier nonmetals that hybridize d
orbitals when they have to.
304
CHAPTER 9
COVALENT BONDING: ORBITALS
305
For sulfur, the valence electrons are in 3s and 3p orbitals. Therefore, 3d orbitals are closest in
energy and are available for hybridization. Arsenic would hybridize 4d orbitals to go with the
valence 4s and 4p orbitals, whereas iodine would hybridize 5d orbitals since the valence
electrons are in n = 5.
12.
Rotation occurs in a bond as long as the orbitals that go to form that bond still overlap when
the atoms are rotating. Sigma bonds, with the head-to-head overlap, remain unaffected by
rotating the atoms in the bonds. Atoms that are bonded together by only a sigma bond (single
bond) exhibit this rotation phenomenon. The π bonds, however, cannot be rotated. The p
orbitals must be parallel to each other to form the π bond. If we try to rotate the atoms in a π
bond, the p orbitals would no longer have the correct alignment necessary to overlap.
Because π bonds are present in double and triple bonds (a double bond is composed of 1 σ
and 1 π bond, and a triple bond is always 1 σ and 2 π bonds), the atoms in a double or triple
bond cannot rotate (unless the bond is broken).
13.
Bonding and antibonding molecular orbitals are both solutions to the quantum mechanical
treatment of the molecule. Bonding orbitals form when in-phase orbitals combine to give
constructive interference. This results in enhanced electron probability located between the
two nuclei. The end result is that a bonding MO is lower in energy than the atomic orbitals
from which it is composed. Antibonding orbitals form when out-of-phase orbitals combine.
The mismatched phases produce destructive interference leading to a node of electron
probability between the two nuclei. With electron distribution pushed to the outside, the
energy of an antibonding orbital is higher than the energy of the atomic orbitals from which it
is composed.
14.
From experiment, B2 is paramagnetic. If the σ2p MO is lower in energy than the two degenerate π2p MOs, the electron configuration for B2 would have all electrons paired. Experiment
tells us we must have unpaired electrons. Therefore, the MO diagram is modified to have the
π2p orbitals lower in energy than the σ2p orbitals. This gives two unpaired electrons in the
electron configuration for B2, which explains the paramagnetic properties of B2. The model
allowed for s and p orbitals to mix, which shifted the energy of the σ2p orbital to above that of
the π2p orbitals.
15.
The localized electron model does not deal effectively with molecules containing unpaired
electrons. We can draw all of the possible structures for NO with its odd number of valence
electrons but still not have a good feel for whether the bond in NO is weaker or stronger than
the bond in NO−. MO theory can handle odd electron species without any modifications.
From the MO electron configurations, the bond order is 2.5 for NO and 2 for NO−. Therefore,
NO should have the stronger bond (and it does). In addition, hybrid orbital theory does not
predict that NO− is paramagnetic. The MO theory correctly makes this prediction.
16.
NO3−, 5 + 3(6) + 1 = 24 e−
O
O
O
N
N
O
O
O
N
O
O
O
306
CHAPTER 9
COVALENT BONDING: ORBITALS
When resonance structures can be drawn, it is usually due to a multiple bond that can be in
different positions. This is the case for NO3−. Experiment tells us that the three N−O bonds
are equivalent in length and strength. To explain this, we say the π electrons are delocalized
in the molecule. For NO3−, the π bonding system is composed of an unhybridized p atomic
orbital from all the atoms in NO3−. These p orbitals are oriented perpendicular to the plane of
the atoms in NO3−. The π bonding system consists of all of the perpendicular p orbitals
overlapping forming a diffuse electron cloud above and below the entire surface of the NO3−
ion. Instead of having the π electrons situated above and below two specific nuclei, we think
of the π electrons in NO3− as extending over the entire surface of the molecule (hence the
term delocalized). See Figure 9.48 for an illustration of the π bonding system in NO3−.
Exercises
The Localized Electron Model and Hybrid Orbitals
17.
H2O has 2(1) + 6 = 8 valence electrons.
O
H
H
H2O has a tetrahedral arrangement of the electron pairs about the O atom that requires sp3
hybridization. Two of the four sp3 hybrid orbitals are used to form bonds to the two
hydrogen atoms, and the other two sp3 hybrid orbitals hold the two lone pairs on oxygen. The
two O−H bonds are formed from overlap of the sp3 hybrid orbitals from oxygen with the 1s
atomic orbitals from the hydrogen atoms. Each O‒H covalent bond is called a sigma (σ)
bond since the shared electron pair in each bond is centered in an area on a line running
between the two atoms.
18.
CCl4 has 4 + 4(7) = 32 valence electrons.
Cl
C
Cl
Cl
Cl
CCl4 has a tetrahedral arrangement of the electron pairs about the carbon atom that requires
sp3 hybridization. The four sp3 hybrid orbitals from carbon are used to form the four bonds to
chlorine. The chlorine atoms also have a tetrahedral arrangement of electron pairs, and we
will assume that they are also sp3 hybridized. The C‒Cl sigma bonds are all formed from
overlap of sp3 hybrid orbitals from carbon with sp3 hybrid orbitals from each chlorine atom.
CHAPTER 9
19.
COVALENT BONDING: ORBITALS
307
H2CO has 2(1) + 4 + 6 = 12 valence electrons.
O
C
H
H
The central carbon atom has a trigonal planar arrangement of the electron pairs that requires
sp2 hybridization. The two C−H sigma bonds are formed from overlap of the sp2 hybrid
orbitals from carbon with the hydrogen 1s atomic orbitals. The double bond between carbon
and oxygen consists of one σ and one π bond. The oxygen atom, like the carbon atom, also
has a trigonal planar arrangement of the electrons that requires sp2 hybridization. The σ bond
in the double bond is formed from overlap of a carbon sp2 hybrid orbital with an oxygen sp2
hybrid orbital. The π bond in the double bond is formed from overlap of the unhybridized p
atomic orbitals. Carbon and oxygen each has one unhybridized p atomic orbital that is
parallel with the other. When two parallel p atomic orbitals overlap, a π bond results where
the shared electron pair occupies the space above and below a line joining the atoms in the
bond.
20.
C2H2 has 2(4) + 2(1) = 10 valence electrons.
H
C
C
H
Each carbon atom in C2H2 is sp hybridized since each carbon atom is surrounded by two
effective pairs of electrons; i.e., each carbon atom has a linear arrangement of the electrons.
Since each carbon atom is sp hybridized, then each carbon atom has two unhybridized p
atomic orbitals. The two C−H sigma bonds are formed from overlap of carbon sp hybrid
orbitals with hydrogen 1s atomic orbitals. The triple bond is composed of one σ bond and
two π bonds. The sigma bond between to the carbon atoms is formed from overlap of sp
hybrid orbitals from each carbon atom. The two π bonds of the triple bond are formed from
parallel overlap of the two unhybridized p atomic orbitals from each carbon.
21.
Ethane, C2H6, has 2(4) + 6(1) = 14 valence electrons.
H
H
H
C
H
C
H
H
The carbon atoms are sp3 hybridized. The six C‒H sigma bonds are formed from overlap of
the sp3 hybrid orbitals from C with the 1s atomic orbitals from the hydrogen atoms. The
carbon-carbon sigma bond is formed from overlap of an sp3 hybrid orbital from each C atom.
Ethanol, C2H6O has 2(4) + 6(1) + 6 = 20 e−
H
H
H
C
H
O
C
H
H
308
CHAPTER 9
COVALENT BONDING: ORBITALS
The two C atoms and the O atom are sp3 hybridized. All bonds are formed from overlap with
these sp3 hybrid orbitals. The C‒H and O‒H sigma bonds are formed from overlap of sp3 hybrid orbitals with hydrogen 1s atomic orbitals. The C‒C and C‒O sigma bonds are formed
from overlap of the sp3 hybrid orbitals from each atom.
22.
HCN, 1 + 4 + 5 = 10 valence electrons
H
N
C
Assuming N is hybridized, both C and N atoms are sp hybridized. The C‒H σ bond is formed
from overlap of a carbon sp3 hybrid orbital with a hydrogen 1s atomic orbital. The triple bond
is composed of one σ bond and two π bonds. The sigma bond is formed from head-to-head
overlap of the sp hybrid orbitals from the C and N atoms. The two π bonds in the triple bond
are formed from overlap of the two unhybridized p atomic orbitals from each C and N atom.
COCl2, 4 + 6 + 2(7) = 24 valence electrons
O
C
Cl
Cl
Assuming all atoms are hybridized, the carbon and oxygen atoms are sp2 hybridized, and the
two chlorine atoms are sp3 hybridized. The two C‒Cl σ bonds are formed from overlap of sp2
hybrids from C with sp3 hybrid orbitals from Cl. The double bond between the carbon and
oxygen atoms consists of one σ and one π bond. The σ bond in the double bond is formed
from head-to-head overlap of an sp2 orbital from carbon with an sp2 hybrid orbital from oxygen. The π bond is formed from parallel overlap of the unhybridized p atomic orbitals from
each atom of C and O.
23.
See Exercises 8.83 and 8.89 for the Lewis structures. To predict the hybridization, first
determine the arrangement of electron pairs about each central atom using the VSEPR model;
then use the information in Figure 9.24 of the text to deduce the hybridization required for
that arrangement of electron pairs.
8.83
8.89
a. CCl4: C is sp3 hybridized.
b. NCl3: N is sp3 hybridized.
c. SeCl2: Se is sp3 hybridized.
d. ICl: Both I and Cl are sp3 hybridized.
a. The central N atom is sp2 hybridized in NO2- and NO3-. In N2O4, both central N
atoms are sp2 hybridized.
b. In OCN- and SCN-, the central carbon atoms in each ion are sp hybridized, and in
N3-, the central N atom is also sp hybridized.
24.
See Exercises 8.84 and 8.90 for the Lewis structures.
8.84
a. All the central atoms are sp3 hybridized.
CHAPTER 9
COVALENT BONDING: ORBITALS
309
b. All the central atoms are sp3 hybridized.
c. All the central atoms are sp3 hybridized.
8.90
In O3 and in SO2, the central atoms are sp2 hybridized, and in SO3, the central sulfur
atom is also sp2 hybridized.
25.
All exhibit dsp3 hybridization. All of these molecules/ions have a trigonal bipyramid
arrangement of electron pairs about the central atom; all have central atoms with dsp3
hybridization. See Exercise 8.87 for the Lewis structures.
26.
All these molecules have an octahedral arrangement of electron pairs about the central atom;
all have central atoms with d2sp3 hybridization. See Exercise 8.88 for the Lewis structures.
27.
The molecules in Exercise 8.113 all have a trigonal planar arrangement of electron pairs
about the central atom, so all have central atoms with sp2 hybridization. The molecules in
Exercise 8.114 all have a tetrahedral arrangement of electron pairs about the central atom, so
all have central atoms with sp3 hybridization. See Exercises 8.113 and 8.114 for the Lewis
structures.
28.
The molecules in Exercise 8.115 all have central atoms with dsp3 hybridization because all
are based on the trigonal bipyramid arrangement of electron pairs. The molecules in Exercise
8.116 all have central atoms with d2sp3 hybridization because all are based on the octahedral
arrangement of electron pairs. See Exercises 8.115 and 8.116 for the Lewis structures.
29.
a.
b.
F
N
F
F
C
F
F
F
F
sp3
nonpolar
tetrahedral
109.5°
trigonal pyramid
<109.5°
sp3
polar
The angles in NF3 should be slightly less than 109.5° because the lone pair requires more
space than the bonding pairs.
c.
d.
O
F
F
F
B
F
V-shaped
<109°.5
3
sp
polar
trigonal planar
120°
F
sp2
nonpolar
310
CHAPTER 9
COVALENT BONDING: ORBITALS
f.
e.
F
H
Be
F b
H
a
Te
F
F
linear
180°
sp
nonpolar
dsp3
polar
h.
g.
F
F
a
As
b
F
F
Kr
F
F
F
dsp3
nonpolar
trigonal bipyramid
a) 90°, b) 120°
i.
see-saw
a) ≈120°, b) ≈90°
F
F
F
j.
F
Kr
dsp3
nonpolar
linear
180°
F
90o
F
Se
F
F
F
F
d2sp3
nonpolar
square planar
90°
d2sp3
nonpolar
octahedral
90°
l.
k.
F
F
F
F
square pyramid
≈90°
I
F
I
F
F
d2sp3
polar
T-shaped
≈90°
F
dsp3
polar
CHAPTER 9
30.
COVALENT BONDING: ORBITALS
311
a.
S
O
V-shaped, sp2, 120°
O
Only one resonance form is shown. Resonance does not change the position of the atoms.
We can predict the geometry and hybridization from any one of the resonance structures.
b.
c.
O
2-
O
S
O
S
O
S
O
O
trigonal planar, 120°, sp2
(plus two other resonance structures)
tetrahedral, 109.5°, sp3
d.
2-
O
O
S
e.
O
O
O
O
S
O
O
O
O
O
trigonal pyramid, <109.5°, sp3
tetrahedral geometry about each S, 109.5°,
sp3 hybrids; V-shaped arrangement about
peroxide O’s, ≈109.5°, sp3
f.
2-
O
g.
S
S
O
2-
S
O
F
F
O
tetrahedral, 109.5°, sp3
V-shaped, <109.5°, sp3
312
CHAPTER 9
COVALENT BONDING: ORBITALS
h.
i.
≈ 90
o
F
F
F
S
o
≈ 120
F
F
S
F
F
F
F
F
see-saw, ≈90° and ≈120°, dsp3
octahedral, 90°, d2sp3
j.
k.
F
b
S
F
c
a
F
F
S
120
F
o
F
a) ≈109.5°
b) ≈90° c) ≈120°
see-saw about S atom with one lone pair (dsp3);
bent about S atom with two lone pairs (sp3)
F
S
+
90
o
F
F
trigonal bipyramid,
90° and 120°, dsp3
31.
H
H
C
H
C
H
For the p orbitals to properly line up to form the π bond, all six atoms are forced into the same
plane. If the atoms are not in the same plane, then the π bond could not form since the p orbitals would no longer be parallel to each other.
32.
No, the CH2 planes are mutually perpendicular to each other. The center C atom is sp
hybridized and is involved in two π bonds. The p orbitals used to form each π bond must be
perpendicular to each other. This forces the two CH2 planes to be perpendicular.
H
H
C
H
C
C
H
CHAPTER 9
33.
COVALENT BONDING: ORBITALS
313
a. There are 33 σ and 9 π bonds. Single bonds always are σ bonds, double bonds always
consist of 1 σ and 1 π bond, and triple bonds always consist of 1 σ and 2 π bonds. The 9
π bonds come from the 9 double bonds in the indigo molecule.
b. All carbon atoms are sp2 hybridized because all have a trigonal planar arrangement of
electron pairs.
34.
The two nitrogen atoms in urea both have a tetrahedral arrangement of electron pairs, so both
of these atoms are sp3 hybridized. The carbon atom has a trigonal planar arrangement of
electron pairs, so C is sp2 hybridized. O is also sp2 hybridized because it also has a trigonal
planar arrangement of electron pairs.
Each of the four N−H sigma bonds are formed from overlap of an sp3 hybrid orbital from
nitrogen with a 1s orbital from hydrogen. Each of the two N−C sigma bonds are formed from
an sp3 hybrid orbital from N with an sp2 hybrid orbital from carbon. The double bond
between carbon and oxygen consists of one σ and one π bond. The σ bond in the double
bond is formed from overlap of a carbon sp2 hybrid orbital with an oxygen sp2 hybrid orbital.
The π bond in the double bond is formed from overlap of the unhybridized p atomic orbitals.
Carbon and oxygen each have one unhybridized p atomic orbital, and they are assumed to be
parallel to each other. When two parallel p atomic orbitals overlap side to side, a π bond
results.
35.
To complete the Lewis structures, just add lone pairs of electrons to satisfy the octet rule for
the atoms with fewer than eight electrons.
Biacetyl (C4H6O2) has 4(4) + 6(1) + 2(6) = 34 valence electrons.
O
H
H
All CCO angles are 120°. The six atoms are not forced to lie in
the same plane because of free rotation about the carboncarbon single (sigma) bonds. There are 11 σ and 2 π bonds in
biacetyl.
O
C
C
H
sp2
sp3
H
C
C
H
H
Acetoin (C4H8O2) has 4(4) + 8(1) + 2(6) = 36 valence electrons.
sp2
sp3
H
O
a
H
C
H
C
H
C
C
b
H
O
H
The carbon with the doubly bonded O is sp2 hybridized.
The other three C atoms are sp3 hybridized. Angle a = 120°
and angle b = 109.5°. There are 13 σ and 1 π bonds in
acetoin.
H
H
sp3
Note: All single bonds are σ bonds, all double bonds are one σ and one π bond, and all triple
bonds are one σ and two π bonds.
314
36.
CHAPTER 9
COVALENT BONDING: ORBITALS
Acrylonitrile: C3H3N has 3(4) + 3(1) + 5 = 20 valence electrons.
H
a) 120°
C
a
b) 120°
H
b
H
C
c
sp2
c) 180°
6 σ and 3 π bonds
C
N
sp
All atoms of acrylonitrile must lie in the same plane. The π bond in the double bond dictates
that the C and H atoms are all in the same plane, and the triple bond dictates that N is in the
same plane with the other atoms.
Methyl methacrylate (C5H8O2) has 5(4) + 8(1) + 2(6) = 40 valence electrons.
d) 120°
H
e) 120°
≈109.5°
C
H
d
sp3
C
O
e
sp2
37.
H
C
C
f)
H
H
O
14 σ and 2 π bonds
H
C
H
f
H
a. Add lone pairs to complete octets for each O and N.
H
a
H
N
C
O
N
b
e
c
N
O
C
d
NH2
H 2C
N
C
C f
g
C
O
O
Azodicarbonamide
CH 3
h
methyl cyanoacrylate
Note: NH2, CH2 (H2C), and CH3 are shorthand for nitrogen or carbon atoms singly
bonded to hydrogen atoms.
b. In azodicarbonamide, the two carbon atoms are sp2 hybridized, the two nitrogen atoms
with hydrogens attached are sp3 hybridized, and the other two nitrogens are sp2
hybridized. In methyl cyanoacrylate, the CH3 carbon is sp3 hybridized, the carbon with
the triple bond is sp hybridized, and the other three carbons are sp2 hybridized.
c. Azodicarbonamide contains three π bonds and methyl cyanoacrylate contains four π
bonds.
CHAPTER 9
COVALENT BONDING: ORBITALS
d. a) ≈109.5°
f) 120°
38.
b) 120°
c) ≈120°
g) ≈109.5°
h) 120°
315
d) 120°
e) 180°
a. Piperine and capsaicin are molecules classified as organic compounds, i.e., compounds
based on carbon. The majority of Lewis structures for organic compounds have all atoms
with zero formal charge. Therefore, carbon atoms in organic compounds will usually
form four bonds, nitrogen atoms will form three bonds and complete the octet with one
lone pair of electrons, and oxygen atoms will form two bonds and complete the octet with
two lone pairs of electrons. Using these guidelines, the Lewis structures are:
Note: The ring structures are all shorthand notation for rings of carbon atoms. In
piperine the first ring contains six carbon atoms and the second ring contains five carbon
atoms (plus nitrogen). Also notice that CH3, CH2, and CH are shorthand for carbon
atoms singly bonded to hydrogen atoms.
b. piperine: 0 sp, 11 sp2 and 6 sp3 carbons; capsaicin: 0 sp, 9 sp2, and 9 sp3 carbons
c. The nitrogens are sp3 hybridized in each molecule.
d. a) 120°
b) 120°
c) 120°
d) 120°
e) ≈109.5°
f) 109.5°
g) 120°
h) 109.5°
i)
120°
j)
k) 120°
l)
109.5°
109.5°
316
39.
CHAPTER 9
COVALENT BONDING: ORBITALS
To complete the Lewis structure, just add lone pairs of electrons to satisfy the octet rule for
the atoms that have fewer than eight electrons.
H
O
C
C
N
H
H
H
O
C
H
H
C
C
H
H
C
N
O
H
O
C
H H C
C
C
N
H
H
N
N
40.
a. 6
b. 4
c. The center N in ‒N=N=N group
d. 33 σ
e. 5 π bonds
f.
g. ≈109.5°
h. sp3
180°
a. To complete the Lewis structure, add two lone pairs to each sulfur atom.
sp3
H 3C
sp
C
sp
C
H
H
C
C
C
C
S
sp
C
sp
C
sp
C
sp
C
sp2
CH
sp2
CH 2
S
b. See the Lewis structure. The four carbon atoms in the ring are all sp2 hybridized, and the
two sulfur atoms are sp3 hybridized.
c. 23 σ and 9 π bonds. Note: CH3 (H3C), CH2, and CH are shorthand for carbon atoms
singly bonded to hydrogen atoms.
CHAPTER 9
COVALENT BONDING: ORBITALS
317
The Molecular Orbital Model
41.
a. The bonding molecular orbital is on the right and the antibonding molecular orbital is on
the left. The bonding MO has the greatest electron probability between the nuclei, while
the antibonding MO has greatest electron probability outside the space between the two
nuclei.
b. The bonding MO is lower in energy. Because the electrons in the bonding MO have the
greatest probability of lying between the two nuclei, these electrons are attracted to two
different nuclei, resulting in a lower energy.
42.
a.
σp
When p orbitals are combined head-to-head and the phases are the same sign (the orbital
lobes have the same sign), a sigma bonding molecular orbital is formed.
b.
πp
When parallel p orbitals are combined in-phase (the signs match up), a pi bonding
molecular orbital is formed.
c.
σp*
When p orbitals are combined head-to-head and the phases are opposite (the orbital lobes
have opposite signs), a sigma antibonding molecular orbital is formed.
d.
πp*
When parallel p orbitals are combined out-of-phase (the orbital lobes have opposite
signs), a pi antibonding molecular orbital is formed.
318
43.
CHAPTER 9
If we calculate a nonzero bond order for a molecule, then we predict that it can exist (is
stable).
a. H2+:
H2:
H2−:
H22−:
(σ1s)1
(σ1s)2
(σ1s)2(σ1s*)1
(σ1s)2(σ1s*)2
B.O. = bond order = (1−0)/2 = 1/2, stable
B.O. = (2−0)/2 = 1, stable
B.O. = (2−1)/2 = 1/2, stable
B.O. = (2−2)/2 = 0, not stable
b. He22+: (σ1s)2
He2+: (σ1s)2(σ1s*)1
He2: (σ1s)2(σ1s*)2
44.
COVALENT BONDING: ORBITALS
B.O. = (2−0)/2 = 1, stable
B.O. = (2−1)/2 = 1/2, stable
B.O. = (2−2)/2 = 0, not stable
B.O. = bond order = (8−4)/2 = 2, stable
a. N22−: (σ2s)2(σ2s*)2(π2p)4(σ2p)2(π2p*)2
2
2
2
4
4
2−
O2 : (σ2s) (σ2s*) (σ2p) (π2p) (π2p*)
B.O. = (8−6)/2 = 1, stable
2
2
2
4
4
2
2−
F2 : (σ2s) (σ2s*) (σ2p) (π2p) (π2p*) (σ2p*) B.O. = (8−8)/2 = 0, not stable
b. Be2: (σ2s)2(σ2s*)2
B.O. = (2−2)/2 = 0, not stable
2
2
2
B2: (σ2s) (σ2s*) (π2p)
B.O. = (4−2)/2 = 1, stable
2
2
2
4
4
2
Ne2: (σ2s) (σ2s*) (σ2p) (π2p) (π2p*) (σ2p*) B.O. = (8−8)/2 = 0, not stable
45.
The electron configurations are:
(σ2s)2
B.O. = (2−0)/2 = 1, diamagnetic (0 unpaired e−)
b. C2:
(σ2s)2(σ2s*)2(π2p)4
B.O. = (6−2)/2 = 2, diamagnetic (0 unpaired e−)
c.
(σ3s)2(σ3s*)2(σ3p)2(π3p)4(π3p*)2
B.O. = (8−4)/2 = 2, paramagnetic (2 unpaired e−)
a.
Li2:
S2:
46.
There are 14 valence electrons in the MO electron configuration. Also, the valence shell is n
= 3. Some possibilities from row 3 having 14 valence electrons are Cl2, SCl−, S22−, and Ar22+.
47.
O2:
(σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)2
B.O. = bond order = (8 – 4)/2 = 2
N2:
(σ2s) (σ2s*) (π2p) (σ2p)
B.O. = (8 – 2)/2 = 3
2
2
4
2
In O2, an antibonding electron is removed, which will increase the bond order to
2.5 [= (8–3)/2]. The bond order increases as an electron is removed, so the bond strengthens.
In N2, a bonding electron is removed, which decreases the bond order to 2.5 = [(7 – 2)/2]. So
the bond strength weakens as an electron is removed from N2.
48.
The electron configurations are:
F2+: (σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)3
B.O. = (8−5)/2 = 1.5; 1 unpaired e−
F2: (σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)4
B.O. = (8−6)/2 = 1;
F2−: (σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)4(σ2p*)1
B.O. = (8−7)/2 = 0.5; 1 unpaired e−
0 unpaired e−
From the calculated bond orders, the order of bond lengths should be F2+ < F2 < F2−.
CHAPTER 9
49.
COVALENT BONDING: ORBITALS
319
For CO, we will assume the same orbital ordering as that for N2.
CO: (σ2s)2(σ2s*)2(π2p)4(σ2p)2
B.O. = (8 – 2)/2 = 3; 0 unpaired electrons
O2: (σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)2
B.O. = (8 – 4)/2 = 2; 2 unpaired electrons
The most obvious differences are that CO has a larger bond order than O2 (3 versus 2) and
that CO is diamagnetic, whereas O2 is paramagnetic.
50.
Considering only the 12 valence electrons in O2, the MO models would be:
σ2p*
↑
↑
↑↓
↑↓
π 2p*
↑↓
π 2p
↑↓
↑↓
↑↓
σ2p
↑↓
↑↓
σ2s*
↑↓
↑↓
σ2s
↑↓
O2 ground state
Arrangement of electrons consistent
with the Lewis structure (double bond
and no unpaired electrons).
It takes energy to pair electrons in the same orbital. Thus the structure with no unpaired
electrons is at a higher energy; it is an excited state.
51.
The electron configurations are (assuming the same orbital order as that for N2):
a. CO: (σ2s)2(σ2s*)2(π2p)4(σ2p)2
B.O. = (8-2)/2 = 3, diamagnetic
b. CO+: (σ2s)2(σ2s*)2(π2p)4(σ2p)1
B.O. = (7-2)/2 = 2.5, paramagnetic
c. CO2+: (σ2s)2(σ2s*)2(π2p)4
B.O. = (6-2)/2 = 2, diamagnetic
Because bond order is directly proportional to bond energy and inversely proportional to
bond length:
Shortest → longest bond length: CO < CO+ < CO2+
Smallest → largest bond energy: CO2+ < CO+ < CO
320
52.
CHAPTER 9
a. CN+:
(σ2s)2(σ2s*)2(π2p)4
b. CN:
(σ2s) (σ2s*) (π2p) (σ2p)
1
(σ2s) (σ2s*) (π2p) (σ2p)
2
−
c. CN :
2
2
2
4
2
4
COVALENT BONDING: ORBITALS
B.O. = (6−2)/2 = 2, diamagnetic
+
B.O. = (7−2)/2 = 2.5, paramagnetic
B.O. = 3, diamagnetic
−
The bond orders are CN , 2; CN, 2.5; CN , 3; because bond order is directly proportional
to bond energy and inversely proportional to bond length:
Shortest → longest bond length: CN− < CN < CN+
Smallest → largest bond energy: CN+ < CN < CN−
53.
a. H2: (σ1s)2
B2: (σ2s)2(σ2s*)2(π2p)2
b. C22−: (σ2s)2(σ2s*)2(π2p)4(σ2p)2
OF: (σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)3
The bond strength will weaken if the electron removed comes from a bonding orbital. Of the
molecules listed, H2, B2, and C22− would be expected to have their bond strength weaken as
an electron is removed. OF has the electron removed from an antibonding orbital, so its bond
strength increases.
54.
a. CN: (σ2s)2(σ2s*)2(π2p)4(σ2p)1
NO: (σ2s)2(σ2s*)2(π2p)4(σ2p)2(π2p*)1
b. O22+: (σ2s)2(σ2s*)2(σ2p)2(π2p)4
N22+: (σ2s)2(σ2s*)2(π2p)4
If the added electron goes into a bonding orbital, the bond order would increase, making the
species more stable and more likely to form. Between CN and NO, CN would most likely
form CN− since the bond order increases (unlike NO− , where the added electron goes into an
antibonding orbital). Between O22+ and N22+, N2+ would most likely form since the bond
order increases (unlike O22+ going to O2+).
55.
The two types of overlap that result in bond formation for p orbitals are in-phase side-to-side
overlap (π bond) and in-phase head-to-head overlap (σ bond).
π2p (in-phase; the signs match up)
σ2p (in-phase; the signs match up)
CHAPTER 9
COVALENT BONDING: ORBITALS
321
56.
+
+
σ* (out-of-phase; the signs
oppose each other)
+
σ
(in-phase; the signs
match up)
These molecular orbitals are sigma MOs because the electron density is cylindrically
symmetric about the internuclear axis.
57.
a. The electron density would be closer to F on average. The F atom is more electronegative than the H atom, and the 2p orbital of F is lower in energy than the 1s orbital of H.
b. The bonding MO would have more fluorine 2p character since it is closer in energy to the
fluorine 2p atomic orbital.
c. The antibonding MO would place more electron density closer to H and would have a
greater contribution from the higher-energy hydrogen 1s atomic orbital.
58.
a. The antibonding MO will have more hydrogen 1s character because the hydrogen 1s
atomic orbital is closer in energy to the antibonding MO.
b. No, the net overall overlap is zero. The px orbital does not have proper symmetry to
overlap with a 1s orbital. The 2px and 2py orbitals are called nonbonding orbitals.
+
+
-
c.
H
HO
O
σ*
↑
1s
↑↓
↑
2p x 2py
↑↓
σ
↑↓
2s
↑
↑↓
↑
2pz
2px
2py
↑↓
2s
d. Bond order = (2 – 0)/2 = 1; Note: The 2s, 2px, and 2py electrons have no effect on the
bond order.
322
CHAPTER 9
COVALENT BONDING: ORBITALS
e. To form OH+, a nonbonding electron is removed from OH. Because the number of
bonding electrons and antibonding electrons is unchanged, the bond order is still equal to
one.
59.
C22− has 10 valence electrons. The Lewis structure predicts sp hybridization for each carbon
with two unhybridized p orbitals on each carbon.
sp hybrid orbitals form the σ bond and the two unhybridized
p atomic orbitals from each carbon form the two π bonds.
2-
C
C
MO: (σ2s)2(σ2s*)2(π2p)4(σ2p)2, B.O. = (8 − 2)/2 = 3
Both give the same picture, a triple bond composed of one σ and two π bonds. Both predict
the ion will be diamagnetic. Lewis structures deal well with diamagnetic (all electrons paired)
species. The Lewis model cannot really predict magnetic properties.
60.
Lewis structures:
NO+:
N
NO:
N
O
O
-
+
NO :
or
N
N
O
O
+ others
Note: Lewis structures do not handle odd numbered electron species very well.
MO model:
NO+: (σ2s)2(σ2s*)2(π2p)4(σ2p)2,
B.O. = 3, 0 unpaired e− (diamagnetic)
NO:
B.O. = 2.5, 1 unpaired e− (paramagnetic)
(σ2s)2(σ2s*)2(π2p)4(σ2p)2(π2p*)1,
NO−: (σ2s)2(σ2s*)2(π2p)4(σ2p)2(π2p*)2
B.O. = 2, 2 unpaired e− (paramagnetic)
The two models give the same results only for NO+ (a triple bond with no unpaired
electrons). Lewis structures are not adequate for NO and NO−. The MO model gives a better
representation for all three species. For NO, Lewis structures are poor for odd electron
species. For NO−, both models predict a double bond, but only the MO model correctly
predicts that NO− is paramagnetic.
61.
O3 and NO2−are isoelectronic, so we only need consider one of them since the same bonding
ideas apply to both. The Lewis structures for O3 are:
O
O
O
O
O
O
For each of the two resonance forms, the central O atom is sp2 hybridized with one unhybridized p atomic orbital. The sp2 hybrid orbitals are used to form the two sigma bonds to the
central atom and hold the lone pair of electrons on the central O atom. The localized electron
view of the π bond uses unhybridized p atomic orbitals. The π bond resonates between the
two positions in the Lewis structures; the actual structure of O3 is an average of the two
resonance structures:
CHAPTER 9
COVALENT BONDING: ORBITALS
323
In the MO picture of the π bond, all three unhybridized p orbitals overlap at the same time,
resulting in π electrons that are delocalized over the entire surface of the molecule. This is
represented as:
or
The Lewis structures for CO32− are (24 e−):
62.
2-
O
C
O
2-
O
C
O
O
2-
O
C
O
O
O
In the localized electron view, the central carbon atom is sp2 hybridized; the sp2 hybrid
orbitals are used to form the three sigma bonds in CO32−. The central C atom also has one
unhybridized p atomic orbital that overlaps with another p atomic orbital from one of the
oxygen atoms to form the π bond in each resonance structure. This localized π bond moves
(resonates) from one position to another. In the molecular orbital model for CO32−, all four
atoms in CO32− have a p atomic orbital that is perpendicular to the plane of the ion. All four
of these p orbitals overlap at the same time to form a delocalized π bonding system where the
π electrons can roam above and below the entire surface of the ion. The π molecular orbital
system for CO32− is analogous to that for NO3− which is shown in Figure 9.48 of the text.
324
CHAPTER 9
COVALENT BONDING: ORBITALS
Additional Exercises
a. XeO3, 8 + 3(6) = 26 e−
63.
b. XeO4, 8 + 4(6) = 32 e−
O
Xe
O
Xe
O
O
O
O
trigonal pyramid; sp3
tetrahedral; sp3
c. XeOF4, 8 + 6 + 4(7) = 42 e−
O
d. XeOF2, 8 + 6 + 2(7) = 28 e−
F
F
F
F
F
Xe
F
O
Xe
or
F
O
F
F
O
Xe
or
F
Xe
O
F
square pyramid; d2sp3
F
T-shaped; dsp3
e. XeO3F2 has 8 + 3(6) + 2(7) = 40 valence electrons.
F
O
F
O
Xe
or
O
Xe
O
O
F
64.
O
F
F
or
Xe
F
trigonal bipyramid; dsp3
O
O
O
FClO2 + F− → F2ClO2−
F3ClO + F− → F4ClO−
F2ClO2−, 2(7) + 7 + 2(6) + 1 = 34 e−
F4ClO−, 4(7) + 7 + 6 + 1 = 42 e−
F
O
-
F
Cl
see-saw, dsp3
F
Cl
O
F
O
F
F
square pyramid, d2sp3
Note: Similar to Exercise 63c, d, and e, F2ClO2− has two additional Lewis structures that are
possible, and F4ClO− has one additional Lewis structure that is possible. The predicted
hybridization is unaffected.
CHAPTER 9
COVALENT BONDING: ORBITALS
325
F3ClO → F− + F2ClO+
F3ClO2 → F− + F2ClO2+
F2ClO+, 2(7) + 7 + 6 − 1 = 26 e−
F2ClO2+, 2(7) + 7 + 2(6) − 1 = 32 e−
+
+
O
Cl
F
F
Cl
O
F
trigonal pyramid, sp3
65.
O
F
tetrahedral, sp3
a. No, some atoms are in different places. Thus these are not resonance structures; they
are different compounds.
b.
For the first Lewis structure, all nitrogen atoms are sp3 hybridized and all carbon atoms
are sp2 hybridized. In the second Lewis structure, all nitrogen atoms and carbon atoms
are sp2 hybridized.
c. For the reaction:
H
O
O
N
C
N
H
H
O
C
C
N
N
C
H
O
O
N
H
C
N
C
O
H
Bonds broken:
3 C=O (745 kJ/mol)
3 C‒N (305 kJ/mol)
3 N‒H (391 kJ/mol)
Bonds formed:
3 C=N (615 kJ/mol)
3 C‒O (358 kJ/mol)
3 O‒H (467 kJ/mol)
ΔH = 3(745) + 3(305) + 3(391) ‒ [3(615) + 3(358) + 3(467)]
ΔH = 4323 kJ ‒ 4320 kJ = 3 kJ
The bonds are slightly stronger in the first structure with the carbon-oxygen double bonds
since ΔH for the reaction is positive. However, the value of ΔH is so small that the best
conclusion is that the bond strengths are comparable in the two structures.
326
66.
CHAPTER 9
COVALENT BONDING: ORBITALS
a. The V-shaped (or bent) molecular structure occurs with both a trigonal planar and a tetrahedral arrangement of electron pairs. If there is a trigonal planar arrangement, the central
atom is sp2 hybridized. If there is a tetrahedral arrangement, the central atom is sp3
hybridized.
b. The see-saw structure is a trigonal bipyramid arrangement of electron pairs which
requires dsp3 hybridization.
c. The trigonal pyramid structure occurs when a central atom has three bonded atoms and a
lone pair of electrons. Whenever a central atom has four effective pairs about the central
atom (exhibits a tetrahedral arrangement of electron pairs), the central atom is sp3
hybridized.
d. A trigonal bipyramidal arrangement of electron pairs requires dsp3 hybridization.
e. A tetrahedral arrangement of electron pairs requires sp3 hybridization.
67.
For carbon, nitrogen, and oxygen atoms to have formal charge values of zero, each C atom
will form four bonds to other atoms and have no lone pairs of electrons, each N atom will
form three bonds to other atoms and have one lone pair of electrons, and each O atom will
form two bonds to other atoms and have two lone pairs of electrons. Following these
bonding requirements gives the following two resonance structures for vitamin B6:
O
b
a
H
O
H
g
H C
H
C
C
H H
c
C
d
e
C
O
C
N
H
H
O
H
C
C
f
O
H
H
H
C
C
C
H
C
H H
C
C
C
O
H
H
C
N
H
a. 21 σ bonds; 4 π bonds (The electrons in the three π bonds in the ring are delocalized.)
b. Angles a), c), and g): ≈109.5°; angles b), d), e), and f): ≈120°
c. 6 sp2 carbons; the five carbon atoms in the ring are sp2 hybridized, as is the carbon with
the double bond to oxygen.
d. 4 sp3 atoms; the two carbons that are not sp2 hybridized are sp3 hybridized, and the
oxygens marked with angles a and c are sp3 hybridized.
e. Yes, the π electrons in the ring are delocalized. The atoms in the ring are all sp2
hybridized. This leaves a p orbital perpendicular to the plane of the ring from each atom.
Overlap of all six of these p orbitals results in a π molecular orbital system where the
electrons are delocalized above and below the plane of the ring (similar to benzene in
Figure 9.47 of the text).
CHAPTER 9
68.
COVALENT BONDING: ORBITALS
327
For carbon, nitrogen, and oxygen atoms to have formal charge values of zero, each C atom
will form four bonds to other atoms and have no lone pairs of electrons, each N atom will
form three bonds to other atoms and have one lone pair of electrons, and each O atom will
form two bonds to other atoms and have two lone pairs of electrons. Following these
bonding requirements, a Lewis structure for aspartame is:
H
H
O
H
H
C
H
O
C
C
H
O
C
H
H
N
C
C
N
C
C
H
H
O
H
H
H
O
H
H
C
C
C
H
C
C
C
H
H
Another resonance structure could be drawn having the double bonds in the benzene ring
moved over one position.
Atoms that have trigonal planar geometry of electron pairs are assumed to have sp2
hybridization, and atoms with tetrahedral geometry of electron pairs are assumed to have sp3
hybridization. All the N atoms have tetrahedral geometry, so they are all sp3 hybridized (no
sp2 hybridization). The oxygens double bonded to carbon atoms are sp2 hybridized; the other
two oxygens with two single bonds are sp3 hybridized. For the carbon atoms, the six carbon
atoms in the benzene ring are sp2 hybridized, and the three carbons double bonded to oxygen
are also sp2 hybridized (tetrahedral geometry). Answering the questions:
•
9 sp2 hybridized C and N atoms (9 from C’s and 0 from N’s)
•
7 sp3 hybridized C and O atoms (5 from C’s and 2 from O’s)
•
39 σ bonds and 6 π bonds (this includes the 3 π bonds in the benzene ring that are
delocalized)
69.
Cl
H
C
H
H
Cl
C
C
Cl
H
Cl
Cl
C
C
Cl
H
C
H
In order to rotate about the double bond, the molecule must go through an intermediate stage
where the π bond is broken and the sigma bond remains intact. Bond energies are 347 kJ/mol
for C‒C and 614 kJ/mol for C=C. If we take the single bond as the strength of the σ bond,
then the strength of the π bond is (614 − 347 = ) 267 kJ/mol. In theory, 267 kJ/mol must be
supplied to rotate about a carbon-carbon double bond.
328
70.
CHAPTER 9
CO, 4 + 6 = 10 e−;
C
CO2, 4 + 2(6) = 16 e−;
O
C
O
COVALENT BONDING: ORBITALS
C3O2, 3(4) + 2(6) = 24 e−
O
O
C
C
C
O
There is no molecular structure for the diatomic CO molecule. The carbon in CO is sp hybridized. CO2 is a linear molecule, and the central carbon atom is sp hybridized. C3O2 is a linear
molecule with all the central carbon atoms exhibiting sp hybridization.
71.
a. BH3 has 3 + 3(1) = 6 valence electrons.
H
trigonal planar, nonpolar, 120°, sp2
B
H
H
b. N2F2 has 2(5) + 2(7) = 24 valence electrons.
Can also be:
N
N
F
F
V-shaped about both N’s;
≈120° about both N’s;
both N atoms: sp2
F
N
N
F
polar
nonpolar
These are distinctly different molecules.
c. C4H6 has 4(4) + 6(1) = 22 valence electrons.
H
H
C
H
C
H
C
C
H
H
All C atoms are trigonal planar with 120° bond angles and sp2 hybridization. Because C
and H have similar electronegativity values, the C−H bonds are essentially nonpolar, so
the molecule is nonpolar. All neutral compounds composed of only C and H atoms are
nonpolar.
72.
a. Yes, both have four sets of electrons about the P. We would predict a tetrahedral
structure for both. See part d for the Lewis structures.
b. The hybridization is sp3 for P in each structure since both structures exhibit a tetrahedral
arrangement of electron pairs.
c. P has to use one of its d orbitals to form the π bond since the p orbitals are all used to
form the hybrid orbitals.
CHAPTER 9
COVALENT BONDING: ORBITALS
329
d. Formal charge = number of valence electrons of an atom − [(number of lone pair
electrons) + 1/2(number of shared electrons)]. The formal charges calculated for the O
and P atoms are next to the atoms in the following Lewis structures.
O -1
+1
Cl
P
0
O
Cl
Cl
P
0
Cl
Cl
Cl
In both structures, the formal charges of the Cl atoms are all zeros. The structure with the
P=O bond is favored on the basis of formal charge since it has a zero formal charge for
all atoms.
73.
a. The Lewis structures for NNO and NON are:
N
N
O
N
N
O
N
N
O
N
O
N
N
O
N
N
O
N
The NNO structure is correct. From the Lewis structures, we would predict both NNO
and NON to be linear. However, we would predict NNO to be polar and NON to be
nonpolar. Since experiments show N2O to be polar, NNO is the correct structure.
b. Formal charge = number of valence electrons of atoms − [(number of lone pair electrons)
+ 1/2(number of shared electrons)].
N
N
O
-1
+1
0
N
N
O
N
N
O
0
+1
-1
-2
+1
+1
The formal charges for the atoms in the various resonance structures are below each
atom. The central N is sp hybridized in all the resonance structures. We can probably
ignore the third resonance structure on the basis of the relatively large formal charges as
compared to the first two resonance structures.
c. The sp hybrid orbitals from the center N overlap with atomic orbitals (or appropriate
hybrid orbitals) from the other two atoms to form the two sigma bonds. The remaining
two unhybridized p orbitals from the center N overlap with two p orbitals from the
peripheral N to form the two π bonds.
2px
sp
sp
2py
z axis
330
74.
CHAPTER 9
COVALENT BONDING: ORBITALS
N2 (ground state): (σ2s)2(σ2s*)2(π2p)4(σ2p)2; B.O. = 3; diamagnetic (0 unpaired e−)
N2 (1st excited state): (σ2s)2(σ2s*)2(π2p)4(σ2p)1(π2p*)1
B.O. = (7 ‒ 3)/2 = 2; paramagnetic (2 unpaired e−)
The first excited state of N2 should have a weaker bond and should be paramagnetic.
75.
F2: (σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)4; F2 should have a lower ionization energy than F. The
electron removed from F2 is in a π2p* antibonding molecular orbital that is higher in energy
than the 2p atomic orbitals from which the electron in atomic fluorine is removed. Because
the electron removed from F2 is higher in energy than the electron removed from F, it should
be easier to remove an electron from F2 than from F.
76.
dxz
+
pz
x
x
z
z
The two orbitals will overlap side to side, so when the orbitals are in phase, a π bonding
molecular orbital would be expected to form.
77.
Side-to-side in-phase overlap of these d orbitals would produce a π bonding molecular
orbital. There would be no probability of finding an electron on the axis joining the two
nuclei, which is characteristic of π MOs.
78.
Molecule A has a tetrahedral arrangement of electron pairs because it is sp3 hybridized.
Molecule B has 6 electron pairs about the central atom, so it is d2sp3 hybridized. Molecule C
has two σ and two π bonds to the central atom, so it either has two double bonds to the central
atom (as in CO2) or one triple bond and one single bond (as in HCN). Molecule C is
consistent with a linear arrangement of electron pairs exhibiting sp hybridization. There are
many correct possibilities for each molecule; an example of each is:
Molecule A: CH4
Molecule B: XeF 4
H
F
F
O
C
O
H
C
N
Xe
C
H
Molecule C: CO2 or HCN
H
F
F
H
tetrahedral; 109.5°; sp3
square planar; 90°; d2sp3
linear; 180°; sp
CHAPTER 9
COVALENT BONDING: ORBITALS
331
ChemWork Problems
The answers to the problems 79-86 (or a variation to these problems) are found in OWL. These
problems are also assignable in OWL.
Challenge Problems
87.
The following Lewis structure has a formal charge of zero for all of the atoms in the
molecule.
H
O
H
H
H
C
C
C
H
H
N
N
C
C
C
C
N
N
O
H
C
H
H
H
The three C atoms each bonded to three H atoms are sp3 hybridized (tetrahedral geometry);
the other five C atoms with trigonal planar geometry are sp2 hybridized. The one N atom with
the double bond is sp2 hybridized, and the other three N atoms are sp3 hybridized. The
answers to the questions are:
88.
•
6 total C and N atoms are sp2 hybridized
•
6 total C and N atoms are sp3 hybridized
•
0 C and N atoms are sp hybridized (linear geometry)
•
25 σ bonds and 4 π bonds
The complete Lewis structure follows on the next page. All but two of the carbon atoms are
sp3 hybridized. The two carbon atoms that contain the double bond are sp2 hybridized (see *
in the following Lewis structure).
332
CHAPTER 9
COVALENT BONDING: ORBITALS
CH3
H
H
C
H2C
H
H
C
CH3
H2C
C
C
*C
H
CH2
CH
CH2
CH3
CH2
CH
H
CH3
C
CH3
C
CH2
C
CH2
H
C
C
H
H
*
C
C
HO
H
H
CH2
H
No; most of the carbons are not in the same plane since a majority of carbon atoms exhibit a
tetrahedral structure (109.5° bond angles). Note: HO, CH, CH2, H2C, and CH3 are shorthand
for oxygen and carbon atoms singly bonded to hydrogen atoms.
89.
a. NCN2− has 5 + 4 + 5 + 2 = 16 valence electrons.
N
N
C
2-
N
N
C
2-
N
C
N
H2NCN has 2(1) + 5 + 4 + 5 = 16 valence electrons.
H +1
N
H
H
0
-1
C
N
0
H
N
0
0
C
N
favored by formal charge
NCNC(NH2)2 has 5 + 4 + 5 + 4 + 2(5) + 4(1) = 32 valence electrons.
H
0
-1
0
+1
N
C
N
0
C
0
0
0
H
N H
N
C
N
0
H
H
0
N
0
C
N
H
N H
0
H
favored by formal charge
2-
CHAPTER 9
COVALENT BONDING: ORBITALS
333
Melamine (C3N6H6) has 3(4) + 6(5) + 6(1) = 48 valence electrons.
H
H
H
N
N
N
C
C
N
N
H
H
N
N
C
H
C
N
H
H
N
C
C
H
N
H
N
N
H
H
b. NCN2−: C is sp hybridized. Each resonance structure predicts a different hybridization for
the N atom. Depending on the resonance form, N is predicted to be sp, sp2, or sp3 hybridized. For the remaining compounds, we will give hybrids for the favored resonance
structures as predicted from formal charge considerations.
NH2
H
N
H
C
N
N
C
sp3
C
N
NH2
sp
2
sp
sp
sp3
Melamine: N in NH2 groups are all sp3 hybridized; atoms in ring are all sp2 hybridized.
c. NCN2−: 2 σ and 2 π bonds; H2NCN: 4 σ and 2 π bonds; dicyandiamide: 9 σ and 3 π
bonds; melamine: 15 σ and 3 π bonds
d. The π-system forces the ring to be planar, just as the benzene ring is planar (see Figure
9.47 of the text).
e. The structure:
N
C
N
C
N
H
H
N
H
H
best agrees with experiments because it has three different CN bonds. This structure is
also favored on the basis of formal charge.
334
90.
CHAPTER 9
COVALENT BONDING: ORBITALS
One of the resonance structures for benzene is:
H
C
H
H
H
C
C
C
C
C
H
H
To break C6H6(g) into C(g) and H(g) requires the breaking of 6 C‒H bonds, 3 C=C bonds,
and 3 C‒C bonds:
C6H6(g) → 6 C(g) + 6 H(g)
ΔH = 6DC‒H + 3DC=C + 3DC‒C
ΔH = 6(413 kJ) + 3(614 kJ) + 3(347 kJ) = 5361 kJ
The question asks for ∆H °f for C6H6(g), which is ΔH for the reaction:
6 C(s) + 3 H2(g) → C6H6(g)
ΔH = ΔH °f , C H ( g )
6 6
To calculate ΔH for this reaction, we will use Hess’s law along with the value ΔH °f for C(g)
and the bond energy value for H2 ( D H 2 = 432 kJ/mol).
6 C(g) + 6 H(g) → C6H6(g)
6 C(s) → 6 C(g)
3 H2(g) → 6 H(g)
ΔH1 = −5361 kJ
ΔH2 = 6(717 kJ)
ΔH3 = 3(432 kJ)
6 C(s) + 3 H2(g) → C6H6(g)
ΔH = ΔH1 + ΔH2 + ΔH3 = 237 kJ; ΔH °f , C H ( g ) = 237 kJ/mol
6 6
The experimental ΔH of for C6H6(g) is more stable (lower in energy) by 154 kJ than the ΔH of
calculated from bond energies (83 − 237 = −154 kJ). This extra stability is related to
benzene’s ability to exhibit resonance. Two equivalent Lewis structures can be drawn for
benzene. The π bonding system implied by each Lewis structure consists of three localized π
bonds. This is not correct because all C‒C bonds in benzene are equivalent. We say the π
electrons in benzene are delocalized over the entire surface of C6H6 (see Section 9.5 of the
text). The large discrepancy between ΔH of values is due to the delocalized π electrons, whose
effect was not accounted for in the calculated ΔH of value. The extra stability associated with
benzene can be called resonance stabilization. In general, molecules that exhibit resonance
are usually more stable than predicted using bond energies.
CHAPTER 9
91.
COVALENT BONDING: ORBITALS
a. E =
335
hc
(6.626 × 10 −34 J s)(2.998 × 108 m / s)
= 7.9 × 10−18 J
=
−9
λ
25 × 10 m
7.9 × 10−18 J ×
6.022 × 10 23
1 kJ
= 4800 kJ/mol
×
mol
1000 J
Using ΔH values from the various reactions, 25-nm light has sufficient energy to ionize
N2 and N and to break the triple bond. Thus N2, N2+, N, and N+ will all be present,
assuming excess N2.
b. To produce atomic nitrogen but no ions, the range of energies of the light must be from
941 kJ/mol to just below 1402 kJ/mol.
1000 J
1 mol
941 kJ
×
×
= 1.56 × 10−18 J/photon
23
1 kJ
mol
6.022 × 10
λ =
(6.6261 × 10 −34 J s)(2.998 × 108 m / s)
hc
= 1.27 × 10−7 m = 127 nm
=
−18
E
1.56 × 10 J
1402 kJ
1 mol
1000 J
×
×
= 2.328 × 10−18 J/photon
23
mol
kJ
6.0221 × 10
λ =
hc
(6.6261 × 10 −34 J s)(2.9979 × 108 m / s)
= 8.533 × 10−8 m = 85.33 nm
=
−18
E
2.328 × 10 J
Light with wavelengths in the range of 85.33 nm < λ < 127 nm will produce N but no
ions.
c. N2: (σ2s)2(σ2s*)2(π2p)4(σ2p)2; the electron removed from N2 is in the σ2p molecular orbital,
which is lower in energy than the 2p atomic orbital from which the electron in atomic
nitrogen is removed. Because the electron removed from N2 is lower in energy than the
electron removed from N, the ionization energy of N2 is greater than that for N.
92.
The π bonds between S atoms and between C and S atoms are not as strong. The p atomic
orbitals do not overlap with each other as well as the smaller p atomic orbitals of C and O
overlap.
93.
O=N‒Cl: The bond order of the NO bond in NOCl is 2 (a double bond).
NO: From molecular orbital theory, the bond order of this NO bond is 2.5. (See Figure 9.40
of the text. )
Both reactions apparently involve only the breaking of the N‒Cl bond. However, in the
reaction ONCl → NO + Cl, some energy is released in forming the stronger NO bond,
lowering the value of ΔH. Therefore, the apparent N‒Cl bond energy is artificially low for
this reaction. The first reaction involves only the breaking of the N‒Cl bond.
336
94.
CHAPTER 9
COVALENT BONDING: ORBITALS
The molecular orbitals for BeH2 are formed from the two hydrogen 1s orbitals and the 2s and
one of the 2p orbitals from beryllium. One of the sigma bonding orbitals forms from overlap
of the hydrogen 1s orbitals with a 2s orbital from beryllium. Assuming the z axis is the
internuclear axis in the linear BeH2 molecule, then the 2pz orbital from beryllium has proper
symmetry to overlap with the 1s orbitals from hydrogen; the 2px and 2py orbitals are
nonbonding orbitals since they don’t have proper symmetry necessary to overlap with 1s
orbitals. The type of bond formed from the 2pz and 1s orbitals is a sigma bond since the
orbitals overlap head to head. The MO diagram for BeH2 is:
Be
2H
σ*s
σ*p
2px
2p
↑↓
2s
↑↓
2py
↑
1s
↑
1s
σp
↑↓
σs
Bond order = (4 − 0)/2 = 2; the MO diagram predicts BeH2 to be a stable species and also
predicts that BeH2 is diamagnetic. Note: The σs MO is a mixture of the two hydrogen 1s
orbitals with the 2s orbital from beryllium, and the σp MO is a mixture of the two hydrogen
1s orbitals with the 2pz orbital from beryllium. The MOs are not localized between any two
atoms; instead, they extend over the entire surface of the three atoms.
95.
The ground state MO electron configuration for He2 is (σ1s)2(σ1s*)2, giving a bond order of 0.
Therefore, He2 molecules are not predicted to be stable (and are not stable) in the lowestenergy ground state. However, in a high-energy environment, electron(s) from the antibonding orbitals in He2 can be promoted into higher-energy bonding orbitals, thus giving a
nonzero bond order and a “reason” to form. For example, a possible excited-state MO
electron configuration for He2 would be (σ1s)2(σ1s*)1(σ2s)1, giving a bond order of (3 – 1)/2 =
1. Thus excited He2 molecules can form, but they spontaneously break apart as the
electron(s) fall back to the ground state, where the bond order equals zero.
CHAPTER 9
COVALENT BONDING: ORBITALS
337
96.
2p
2s
O2
O2−
O2+
O
The order from lowest IE to highest IE is: O2− < O2 < O2+ < O.
The electrons for O2−, O2, and O2+ that are highest in energy are in the π *2 p MOs. But for O2−,
these electrons are paired. O2− should have the lowest ionization energy (its paired π *2 p
electron is easiest to remove). The species O2+ has an overall positive charge, making it
harder to remove an electron from O2+ than from O2. The highest-energy electrons for O (in
the 2p atomic orbitals) are lower in energy than the π *2 p electrons for the other species; O will
have the highest ionization energy because it requires a larger quantity of energy to remove
an electron from O as compared to the other species.
97.
The electron configurations are:
N2:
(σ2s)2(σ2s*)2(π2p)4(σ2p)2
O2:
(σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)2
Note: The ordering of the σ2p and π2p orbitals
is not important to this question.
N22−: (σ2s)2(σ2s*)2(π2p)4(σ2p)2(π2p*)2
N2−: (σ2s)2(σ2s*)2(π2p)4(σ2p)2(π2p*)1
O2+: (σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)1
The species with the smallest ionization energy has the electron that is easiest to remove.
From the MO electron configurations, O2, N22−, N2−, and O2+ all contain electrons in the same
higher-energy antibonding orbitals ( π *2 p ) , so they should have electrons that are easier to
remove as compared to N2, which has no π *2 p electrons. To differentiate which has the easiest
π *2 p to remove, concentrate on the number of electrons in the orbitals attracted to the number
of protons in the nucleus.
N22− and N2− both have 14 protons in the two nuclei combined. Because N22− has more
electrons, one would expect N22− to have more electron repulsions, which translates into
having an easier electron to remove. Between O2 and O2+, the electron in O2 should be easier
to remove. O2 has one more electron than O2+, and one would expect the fewer electrons in
338
CHAPTER 9
COVALENT BONDING: ORBITALS
O2+ to be better attracted to the nuclei (and harder to remove). Between N22− and O2, both
have 16 electrons; the difference is the number of protons in the nucleus. Because N22− has
two fewer protons than O2, one would expect the N22− to have the easiest electron to remove,
which translates into the smallest ionization energy.
98.
a. F2−(g) → F(g) + F−(g) ∆H = F2− bond energy
Using Hess’s law:
F2−(g) → F2(g) + e−
F2(g) → 2 F(g)
F(g) + e− → F−(g)
F2−(g) → F(g) + F−(g)
∆H = 290. kJ (ionization energy for F2−)
∆H = 154 kJ (bond energy for F2 from Table 8.4)
∆H = –327.8 kJ (electron affinity for F from Table 7.7)
∆H = 116 kJ; bond energy for F2− = 116 kJ/mol
Note that F2− has a smaller bond energy than F2.
b. F2: (σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)4
F2−:
(σ2s) (σ2s*) (σ2p) (π2p) (π2p*) (σ2p*)
2
2
2
4
4
B.O. = (8 – 6)/2 = 1
1
B.O. = (8 – 7)/2 = 0.5
MO theory predicts that F2 should have a stronger bond than F2− because F2 has the larger
bond order. As determined in part a, F2 indeed has a stronger bond because the F2 bond
energy (154 kJ/mol) is greater than the F2− bond energy (116 kJ/mol).
99.
a. The CO bond is polar with the negative end at the more electronegative oxygen atom.
We would expect metal cations to be attracted to and bond to the oxygen end of CO on
the basis of electronegativity.
b.
C
O
FC (carbon) = 4 − 2 − 1/2(6) = −1
FC (oxygen) = 6 − 2 − 1/2(6) = +1
From formal charge, we would expect metal cations to bond to the carbon (with the
negative formal charge).
c. In molecular orbital theory, only orbitals with proper symmetry overlap to form bonding
orbitals. The metals that form bonds to CO are usually transition metals, all of which
have outer electrons in the d orbitals. The only molecular orbitals of CO that have proper
symmetry to overlap with d orbitals are the π2p* orbitals, whose shape is similar to the d
orbitals. Because the antibonding molecular orbitals have more carbon character (carbon
is less electronegative than oxygen), one would expect the bond to form through carbon.
CHAPTER 9
100.
COVALENT BONDING: ORBITALS
339
Benzoic acid (C7H6O2) has 7(4) + 6(1) + 2(6) = 46 valence electrons. The Lewis structure for
benzoic acid is:
O
C
H
H
C
C
C
C
H
H
O
The circle in the ring indicates the delocalized
π bonding in the benzene ring. The two benzene resonance Lewis structures have three
alternating double bonds in the ring (see
Figure 9.45).
C
C
H
H
The six carbons in the ring and the carbon bonded to the ring are all sp2 hybridized. The five
C−H sigma bonds are formed from overlap of the sp2 hybridized carbon atoms with hydrogen
1s atomic orbitals. The seven C−C σ bonds are formed from head to head overlap of sp2
hybrid orbitals from each carbon. The C−O single bond is formed from overlap of an sp2
hybrid orbital on carbon with an sp3 hybrid orbital from oxygen. The C−O σ bond in the
double bond is formed from overlap of carbon sp2 hybrid orbital with an oxygen sp2 orbital.
The π bond in the C−O double bond is formed from overlap of parallel p unhybridized atomic
orbitals from C and O. The delocalized π bonding system in the ring is formed from overlap
of all six unhybridized p atomic orbitals from the six carbon atoms. See Figure 9.47 for
delocalized π bonding system in the benzene ring.
Integrative Problems
101.
a. Li2: (σ 2s ) 2
B.O. = (2 – 0)/2 = 1
B2: (σ 2s ) 2 (σ *2s ) 2 ( π 2 p ) 2
B.O. = (4 – 2)/2 = 1
Both have a bond order of 1.
b. B2 has four more electrons than Li2, so four electrons must be removed from B2 to make
it isoelectronic with Li2. The isoelectronic ion is B24+.
c. To form B24+, it takes 6455 kJ of energy to remove 4 mol of electrons from 1 mol of B2.
1.5 kg B2 ×
1 mol B 2
1000 g
6455 kJ
×
×
= 4.5 × 105 kJ
1 kg
21.62 g B 2
mol B 2
340
102.
CHAPTER 9
a. HF, 1 + 7 = 8 e−
COVALENT BONDING: ORBITALS
SbF5, 5 + 5(7) = 40 e−
F
F
H
F
F
Sb
F
F
linear, sp3 (if F is hybridized)
H2F+, 2(1) + 7 – 1 = 8 e−
+
F
H
H
trigonal bipyramid, dsp3
SbF6−, 5 + 6(7) + 1 = 48 e−
-
F
F
F
Sb
F
F
F
V-shaped, sp3
b. 2.93 mL ×
octahedral, d2sp3
256.8 g
0.975 g HF
1 mol HF
1 mol [H 2 F]+ [SbF6 ]−
×
×
×
mL
20.01 g HF
2 mol HF
mol [H 2 F]+ [SbF6 ]−
= 18.3 g [H2F]+[SbF6]−
10.0 mL ×
256.8 g
3.10 g SbF5
1 mol SbF5
1 mol [H 2 F] + [SbF6 ] −
×
×
×
mL
216.8 g SbF5
mol SbF5
mol [H 2 F] + [SbF6 ] −
= 36.7 g [H2F]+[SbF6]−
Because HF produces the smaller amount of product, HF is limiting and 18.3 g of
[H2F]+[SbF6]− can be produced.
103.
Element X has 36 protons, which identifies it as Kr. Element Y has one less electron than Y−,
so the electron configuration of Y is 1s22s22p5. This is F.
KrF3+, 8 + 3(7) – 1 = 28 e−
F
F
Kr
F
+
T-shaped, dsp3
CHAPTER 10
LIQUIDS AND SOLIDS
Questions
12.
Chalk is composed of the ionic compound calcium carbonate (CaCO3). The electrostatic
forces in ionic compounds are much stronger than the intermolecular forces in covalent
compounds. Therefore, CaCO3 should have a much higher boiling point than the covalent
compounds found in motor oil and in H2O. Motor oil is composed of nonpolar C−C and C−H
bonds. The intermolecular forces in motor oil are therefore London dispersion forces. We
generally consider these forces to be weak. However, with compounds that have large molar
masses, these London dispersion forces add up significantly and can overtake the relatively
strong hydrogen-bonding interactions in water.
13.
Answer a is correct. Intermolecular forces are the forces between molecules that hold the
substances together in the solid and liquid phases. Hydrogen bonding is a specific type of
intermolecular forces. In this figure, the dotted lines represent the hydrogen bonding
interactions that hold individual H2O molecules together in the solid and liquid phases. The
solid lines represent the O−H covalent bonds.
14.
Hydrogen bonding occurs when hydrogen atoms are covalently bonded to highly
electronegative atoms such as oxygen, nitrogen, or fluorine. Because the electronegativity
difference between hydrogen and these highly electronegative atoms is relatively large, the
N−H, O−H, and F−H bonds are very polar covalent bonds. This leads to strong dipole forces.
Also, the small size of the hydrogen atom allows the dipoles to approach each other more
closely than can occur between most polar molecules. Both of these factors make hydrogen
bonding a special type of dipole interaction.
15.
Atoms have an approximately spherical shape (on average). It is impossible to pack spheres
together without some empty space among the spheres.
16.
Critical temperature: The temperature above which a liquid cannot exist; i.e., the gas cannot
be liquified by increased pressure.
Critical pressure: The pressure that must be applied to a substance at its critical temperature
to produce a liquid.
341
342
CHAPTER 10
LIQUIDS AND SOLIDS
The kinetic energy distribution changes as one raises the temperature (T4 > Tc > T3 > T2 >
T1). At the critical temperature Tc, all molecules have kinetic energies greater than the
intermolecular forces F, and a liquid can't form. Note: The various temperature distributions
shown in the plot are not to scale. The area under each temperature distribution should be
equal to each other (area = total number of molecules).
17.
Evaporation takes place when some molecules at the surface of a liquid have enough energy
to break the intermolecular forces holding them in the liquid phase. When a liquid
evaporates, the molecules that escape have high kinetic energies. The average kinetic energy
of the remaining molecules is lower; thus the temperature of the liquid is lower.
18.
A crystalline solid will have the simpler diffraction pattern because a regular, repeating
arrangement is necessary to produce planes of atoms that will diffract the X rays in regular
patterns. An amorphous solid does not have a regular repeating arrangement and will
produce a complicated diffraction pattern.
19.
An alloy is a substance that contains a mixture of elements and has metallic properties. In a
substitutional alloy, some of the host metal atoms are replaced by other metal atoms of
similar size, e.g., brass, pewter, plumber’s solder. An interstitial alloy is formed when some
of the interstices (holes) in the closest packed metal structure are occupied by smaller atoms,
e.g., carbon steels.
20.
Equilibrium: There is no change in composition; the vapor pressure is constant.
Dynamic: Two processes, vapor → liquid and liquid → vapor, are both occurring but with
equal rates, so the composition of the vapor is constant.
21.
a. As the strength of the intermolecular forces increase, the rate of evaporation decreases.
b. As temperature increases, the rate of evaporation increases.
c. As surface area increases, the rate of evaporation increases.
22.
H2O(l) → H2O(g) ∆H° = 44 kJ/mol; the vaporization of water is an endothermic process.
In order to evaporate, water must absorb heat from the surroundings. In this case, part of the
surroundings is our body. So, as water evaporates, our body supplies heat, and as a result,
our body temperature can cool down. From Le Châtelier’s principle, the less water vapor in
the air, the more favorable the evaporation process. Thus the less humid the surroundings,
the more favorably water converts into vapor, and the more heat that is lost by our bodies.
23.
C2H5OH(l) → C2H5OH(g) is an endothermic process. Heat is absorbed when liquid ethanol
vaporizes; the internal heat from the body provides this heat, which results in the cooling of
the body.
24.
The phase change H2O(g) → H2O(l) releases heat that can cause additional damage. Also,
steam can be at a temperature greater than 100°C.
25.
Sublimation will occur, allowing water to escape as H2O(g).
CHAPTER 10
LIQUIDS AND SOLIDS
343
26.
Water boils when the vapor pressure equals the external pressure. Because the external
pressure is significantly lower at high altitudes, a lower temperature is required to equalize
the vapor pressure of water to the external pressure. Thus food cooked in boiling water at
high elevations cooks at a lower temperature, so it takes longer.
27.
The strength of intermolecular forces determines relative boiling points. The types of
intermolecular forces for covalent compounds are London dispersion forces, dipole forces,
and hydrogen bonding. Because the three compounds are assumed to have similar molar mass
and shape, the strength of the London dispersion forces will be about equal among the three
compounds. One of the compounds will be nonpolar, so it only has London dispersion forces.
The other two compounds will be polar, so they have additional dipole forces and will boil at
a higher temperature than the nonpolar compound. One of the polar compounds will have an
H covalently bonded to either N, O, or F. This gives rise to the strongest type of covalent
intermolecular forces, hydrogen bonding. The compound that hydrogen bonds will have the
highest boiling point, whereas the polar compound with no hydrogen bonding will boil at a
temperature in the middle of the other compounds.
28.
a. Both forms of carbon are network solids. In diamond, each carbon atom is surrounded by
a tetrahedral arrangement of other carbon atoms to form a huge molecule. Each carbon
atom is covalently bonded to four other carbon atoms.
The structure of graphite is based on layers of carbon atoms a